ESSENTIALS   OF   CHEMISTRY 


BY 


JOHN    C.   HESSLER,   PH.D. 

t » 

INSTRUCTOR  IN  CHEMISTRY,  THE  UNIVERSITY  OF  CHICAGO;    LATE 

INSTRUCTOR  IN  CHEMISTRY,  THE  HYDE    PARK 

(CHICAGO)    HIGH  SCHOOL 


AND 


ALBERT  L.  SMITH,   PH.D. 

INSTRUCTOR  IN   CHEMISTRY,  THE   ENGLEWOOD  HIGH   SCHOOL, 
CHICAGO 


ov    TroXX'    aXXa 


BOSTON,  U.S.A. 

BENJ.  H.  SANBORN  &  CO. 


COPYRIGHT,  1902, 
By  JOHN  C.  HESSLER  AND  ALBERT  L.  SMITH 


SHje  JFort 

SAMUEL    USHER 

176  TO   184  HIGH  STREET 

BOSTON,  MASS. 


PREFACE. 


The  interest  aroused  by  the  introduction  of  laboratory 
science  into  secondary  schools,  a  decade  and  a  half  ago, 
brought  out  a  number  of  text-books  of  Chemistry.  Al- 
though most  of  these  books  had  directions  for  laboratory 
work  scattered  through  the  descriptive  matter,  yet  the 
laboratory  exercises  Avere  illustrations  of  the  text  proper 
rather  than  working  notes  for  the  student.  A  further 
source  of  confusion  lay  in  the  fact  that  the  experiments 
capable  of  being  performed  by  the  student  were  hope- 
lessly mixed  with  those  intended  for  the  teacher. 
When  the  separate  "  laboratory  manual "  appeared  it  was 
usually  characterized  by  the  same  faults. 

Within  recent  years  many  new  text-books  of  Chem- 
istry have  been  written  for  secondary  schools,  but,  with 
a  few  exceptions,  the  new  books,  like  the  old  ones,  are 
impractical.  They  are  either  too  diffuse  in  description, 
the  laboratory  work  being  left  chiefly  to  the  invention  of 
the  teacher,  or  they  are  merely  laboratory  manuals, 
without  enough  descriptive  matter  to  make  them  useful 
as  text-books.  Moreover,  the  laboratory  exercises  of 
many  modern  books  are  wholly  beyond  the  capabilities  of 
the  average  student. 

There  is,  therefore,  a  demand  for  a  text-book  of 
Chemistry  containing  an  adequate  and  scientific  account 

221752 


iv  PREFACE. 

of  such  of  the  fundamental  facts,  laws,  and  theories  of 
the  subject  as  are  adapted  to  the  needs  of  secondary 
schools,  and  also  specific  directions  for  the  laboratory 
work,  —  directions  that  have  been  tested  and  found  prac- 
ticable. This  demand  it  is  the  design  of  the  present 
book  to  Lieet. 

The  authors  have  had  exceptional  opportunities  to 
know  the  capacity  of  the  average  student  and  the  train- 
ing of  the  average  teacher  of  science  in  secondary 
schools.  They  are  fully  aware  of  the  limitations  of 
these  schools  both  as  to  laboratory  equipment  and  as  to 
the  time  which  may  reasonably  be  expected  for  the  study 
of  Chemistry.  They  have  prepared  the  book  with  these 
limitations  constantly  in  mind. 

A  connected  treatment  of  the  descriptive  matter  of 
this  work  is  attained  by  the  division  of  the  book  into 
three  parts:  (1)  the  text  proper;  (2)  the  laboratory 
exercises  ;  (3)  the  handbook.  The  text  and  the  labora- 
tory exercises  are  bound  together ;  the  handbook  is  in 
pamphlet  form. 

The  text  proper  may  be  characterized  by  saying  that 
it  recognizes  the  fact  that  the  terms  and  the  ideas  of 
Chemistry  are  outside  of  the  common  experience,  and  that 
it  is  useless  to  expect  the  pupil  to  grasp  theoretical  con- 
ceptions before  he  has  become  acquainted  with  the  fun- 
damental phenomena  of  the  science.  The  arrangement 
of  topics  is  such  that  the  early  chapters  of  the  book  are 
mainly  descriptive  ;  theoretical  ideas  are  not  introduced 
until  later,  and  then  only  in  an  elementary  manner. 


PREFACE.  V 

An  illustration  of  the  arrangement  is  the  case  of 
molecular  masses,  atoms,  and  atomic  masses,  which  are 
not  mentioned  at  all  until  Chapter  XVI,  after  the  ele- 
ments hydrogen,  oxygen,  chlorine,  nitrogen,  sulphur, 
and  carbon,  and  their  most  important  compounds,  have 
been  studied. 

Equations,  although  introduced  early  to  accustom  the 
pupil  to  the  fact  that  chemical  reactions  are  quantitative, 
are  for  some  time  (up  to  page  69)  written  out  in  full 
rather  than  in  symbolic  form. 

The  idea  of  equilibrium  is  introduced  first  in  connec- 
tion with  the  study  of  diffusion ;  later  its  application  is 
extended  to  chemical  reactions. 

The  aim  has  been  to  make  the  text  modern  and 
scientific,  yet  not  too  difficult  for  secondary  students. 

The  laboratory  exercises  are  placed  together  after 
the  text.  The  directions  for  these  exercises  are  specific. 
The  quantities  of  reagents  to  be  employed  have  been 
carefully  considered.  TJie  apparatus  required  is  simple 
and  within  the  reach  of  every  school.  The  experiments 
are  so  arranged  that  they  may  be  used  in  schools  in 
which  only  one-hour  laboratory  periods  are  possible,  as 
well  as  in  those  able  to  give  two  or  more  hours  of  con- 
secutive time. 

The  experiments  are  mainly  qualitative,  as  experience 
has  shown  that  they  must  be,  even  with  beginning  college 
students.  Several  experiments  of  a  quantitative  nature 
have,  however,  been  introduced;  for  these  only  the 
simplest  apparatus  is  required. 


Vi  PREFACE. 

In  the  earlier  experiments,  especially,  the  directions 
are  very  explicit;  the  teacher  should,  therefore,  be  able 
to  use  his  laboratory  time  from  the  beginning  in  question- 
ing individual  students  and  in  directing  their  work  at 
close  range. 

Much  material  designed  chiefly  for  the  teacher's  use 
has  been  put  into  the  handbook.  The  handbook  con- 
tains, also,  a  list  of  experiments  to  be  performed  by  the 
teacher  before  the  class.  These  experiments  are  care- 
fully planned,  and  the  directions  for  carrying  them  out, 
accurate.  Every  teacher  will  thus  have  at  his  command 
a  series  of  demonstrations  with  which  to  supplement  the 
laboratory  exercises  performed  by  the  students. 

The  authors  suggest  that  the  several  parts  of  this 
book  be  used  as  follows :  — 

(1)  The  student  is  to  perform  the  experiments  for 
the  week  in  the  laboratory,  taking  notes  upon  his  work. 
These  notes  are  to  be  examined  weekly  by  the  te:  cher. 

(2)  The  teacher  is  to  perform  the  demonstration  ex- 
periments for  the  week,  the  class  taking  notes. 

(3)  The   teacher  is   to  assign   topics  for  recitation. 
These   topics   should   demand   (a)    an   account    of   the 
laboratory  work,  (£>)  a  description  of  the  teacher's  ex- 
periments, (<?)  a  study  of  selected  portions  of  the  text- 
book (supplemented  by  other  books,  if  possible),  and 
(d)  the  working  of  certain  exercises  selected  from  those 
given  at  the  end  of  almost  every  chapter. 

The  authors  believe  that  if  there  are  five  one-hour 
periods  allowed  for  Chemistry,  the  best  results  will  be 


PREFACE.  Vll 

obtained  if  two  of  these  periods  are  given  to  laboratory 
work,  one  to  the  teacher's  experiments,  and  two  to  reci- 
tation. Other  plans  of  work  may,  of  course,  be  better 
adapted  to  individual  teachers. 

The  authors  acknowledge  their  indebtedness  to  the 
long  list  of  persons  who  have  written  upon  elementary 
Chemistry  in  the  past,  and  without  whose  work  the 
preparation  of  a  book  like  this  would  be  almost  im- 
possible. 

They  wish,  further,  to  express  their  thanks  to  Dr. 
Alexander  Smith,  of  the  University  of  Chicago,  who 
suggested  the  apparatus  and  method  of  making  sul- 
phuric acid  (Fig.  41) ;  to  Dr.  Oliver  C.  Farrington,  of 
the  Field  Columbian  Museum,  Chicago,  and  to  Mr.  Ben 
Hains,  Crawfordsville,  Ind.,  for  the  cut  of  Marengo 
Cave  (Fig.  46) ;  to  Mr.  A.  H.  Hutchinson,  manager  of 
the  Ice  and  Refrigerating  Department  of  the  Frick 
Company,  Waynesboro,  Pa.,  for  the  cut  of  the  am- 
monia apparatus  (Fig.  34),  and  to  the  United  Gas 
Improvement  Company,  Philadelphia,  Pa.,  for  the  cut 
of  the  water-gas  apparatus  (Fig.  50). 

Corrections  or  suggestions  from  teachers  using  this 
book,  or  from  other  persons  interested,  will  be  gratefully 
received  by  the  authors. 

J.  C.  H. 

A.  L.  S. 
CHICAGO,  ILL.,  April,  1902. 


PUBLISHER'S   NOTE. 

INSTKUCTOKS  in  Chemistry  need  not  be  told  that  it  is 
now  the  custom  of  lecturers  on  this  subject  to  use  the 
arrow  interchangeably  with,  and  more  often  in  place  of, 
the  equality  sign.  This  is  done  in  this  book.  It  makes 
the  equation  plain,  and,  to  the  chemist,  means  more 
than  does  the  usual  (old)  sign  of  equality.  Its  use  is 
fully  explained  on  page  69.  The  double  arrow,  first 
used  on  page  263,  shows  a  u  reversible  reaction,"  which 
cannot  be  represented  by  the  usual  equality  signs. 


CONTENTS. 


PART  I. 

PAGE. 
Introduction 1 

Definition  of  Chemistry.  —  Importance  of  Chemistry.  — 
Relation  between  Chemistry  and  Physics.  —  Reagents  and 
Reactions.  —  Elements  and  Compounds.  —  Relative  Abun- 
dance and  Importance  of  the  Elements.  —  Exercises. 


CHAPTER   I. 

/^"Hydrogen 9 

Existence.  —  Common  Method  of  Preparation.  —  Purifica- 
tion.—  Chemical  Properties.  —  Union  of  Hydrogen  and  Oxy- 
gen. —  Occlusion.  —  Physical  Properties.  —  Other  Methods  of 
Preparation.  —  Quantitative  Character  of  Chemical  Changes. 
—  Calculation  of  Quantities  of  Factors  and  Products.  —  Ex- 
ercises. 

CHAPTER  II. 

-  Oxygen 24 

Preparation.  —  Common  Laboratory  Method.  —  Other  Meth- 
ods. —  Physical  Properties.  —  Chemical  Properties.  —  Oxides. 

—  Oxidation    and    Reduction.  —  Deflagration.  —  Combustion. 

—  Slow   Combustion. — Spontaneous   Combustion. — Ignition 
Temperature.  —  Combustion  in  Air.  —  Drafts.  —  Safety  Lamp. 

—  Flames.  —  Reversed  Combustion.  —  Exercises. 

ix 


CONTENTS. 
CHAPTER  III. 


PAGE. 

38 


Nature  of  Water.  —  Electrolysis  of  Water.  —  Synthesis  of 
Water  by  Volume  and  by  Weight.  —  Natural  Water.  —  Sea 
Water.  —  Purification  of  Water.  —  Hard  and  Soft  Water.  — 
/  Properties  of  Water.  —  Steam  and  its  Dissociation.  —  Action 
of  Sodium  upon  Water.  —  Hydroxides.  —  Action  of  Metals  upon 
Hydroxides.  —  Water  in  Combination.  —  Water  Mechanically 
Enclosed.  —  Water  of  Crystallization.  —  Water  an  Integral 
Part  of  a  Substance.  —  Efflorescence.  —  Deliquescence.  — 
Exercises. 


CHAPTER  IV. 

^Solution 58 

Character  of  Solution.  —  Boiling  Point  and  Freezing  Point 
of  a  Solution.  —  Temperature  Changes  during  Solution.  — 
Solubility.  —  Effect  of  Temperature  on  Solution.  —  Soluble 
and  Insoluble  Substances.  —  Saturated  Solution.  —  Super- 
saturated Solutions.  —  Precipitation  and  Crystallization.  — 
Decantation  and  Filtration.  —  Crystallization  from  Fusion.  — 
Effervescence.  —  Exercises. 


CHAPTER  V. 

/^Fundamental  Laws,  Combining  Numbers  and  Nomenclature  .  66 
Persistence  of  Mass.  —  Constant  Proportions.  —  Symbols 
and  Formulas.  —  Symbolic  Equations. — Equations  the  Result 
of  Experiment.  —  Quantitative  Meaning  of  Symbols  and  For- 
mulas.—  Multiples  of  Symbols  and  Formulas. — Combining 
Proportions.  —  How  a  Compound  of  Two  Elements  is  Named. 
—  How  to  Distinguish  between  Compounds  of  the  Same  Two 
Elements.  —  Exercises. 


CONTENTS.  XI 

CHAPTER  VI. 

PAGE. 

-Chlorine  ....................     70 

Existence.  —  Common  Method  of  Preparation.  —  Other 
Methods.  —  Physical  Properties.  —  Chemical  Properties.  — 
Action  of  Chlorine  and  Water.  —  Chlorine  and  Ammonia.  — 
Chlorine  and  Turpentine.  —  Uses.  —  Exercises. 


CHAPTER  VII. 

/  Hydrochloric  Acid     ................     88 

Existence.  —  Preparation.  —  Commercial  Manufacture.  — 
Physical  Properties.  —  Volumetric  Composition.  —  Acid  Prop- 
erties. —  Chlorides.  —  Exercises. 


CHAPTER  VIII. 

V^Acids,  Bases,  and  Salts 95 

Acids.  —  Bases.  —  Neutralization.  —  Action  of  Oxides  with 
Acids. — Salts.  —  Acid  Salts.  —  Basic  Salts.  —  Basicity  and 
Acidity.  —  Nomenclature  of  Acids,  Salts,  and  Bases.  —  Exer- 
cises. 


CHAPTER  IX. 

Nitrogen  and  the  Atmosphere 107 

Existence  of  Nitrogen.  —  Preparation.  —  Properties.  —  "  At- 
mospheric Nitrogen"  a  Mixture.  —  Character  of  the  Atmos- 
phere.—  Argon.  —  Helium.  —  Carbon  Dioxide  in  the  Atmos- 
phere. —  Water  Vapor  in  the  Atmosphere.  —  Atmospheric 
Dust.  —  Weight  and  Pressure  of  the  Atmosphere.  —  Lique- 
faction of  Air.  —  Properties  of  Liquid  Air.  —  Determination 
of  the  Proportion  of  Oxygen  in  Air.  —  Air  a  Physical  Mix- 
ture. —  Exercises. 


CONTENTS. 

CHAPTER  X. 

PAGE. 

of  Gases.     The  Molecular  Theory 121 

Gases  and  Vapors.  —  Relation  of  the  Volume  of  a  Gas 
to  Pressure.  —  Relation  of  Volume  to  Temperature.  —  Reduc- 
tion to  Standard  Temperature  and  Pressure.  —  Correction  of 
the  Barometric  Reading.  —  Molecular  Theory.  —  Physical 
States  of  Matter. — Avogadro's  Hypothesis.  —  Explanation  of 
Diffusion.  —  Diffusion  of  Gases  and  of  Liquids.  —  Solution  of 
Gases  in  Liquids.  —  Of  Solids  in  Liquids.  —  Osmotic  Pressure. 
—  Its  Laws.  —  Exercises. 


CHAPTER  XI. 

Ammonia. 136 

Existence.  —  Laboratory  Method  of  Preparation.  —  Com- 
mercial Sources.  —  Physical  Properties.  —  Liquefaction  of 
Ammonia.  —  Liquid  Ammonia  as  a  Refrigerating  Agent.  — 
Chemical  Properties  of  Ammonia.  —  Ammonium  Compounds. 
—  Their  Dissociation.  —  Composition"  of  Ammonia.  —  Synthe- 
sis of  Ammonia  from  Nitrogen  and  Hydrogen.  —  Action  of 
Chlorine  on  Ammonia.  —  Exercises. 


CHAPTER  XII. 

Nitrogen  Acids  and  Oxides 149 

Nitric  Acid.  —  Commercial  Preparation.  —  Laboratory 
Method.  —  Preparation  Compared  with  that  of  Hydrochloric 
Acid.  —  Properties  of  Nitric  Acid. — Action  upon  Metals. — 
Aqua  Regia.  —  Oxidation  by  Nitric  Acid.  —  Explanation.  — 
Formation  of  Nitrates  in  Nature.  —  Manufacture  of  Potas- 
sium Nitrate. — Uses  of  Nitric  Acid  and  the  Nitrates. — 
Nitrogen  Pentoxide. — Nitrous  Acid. — Nitrogen  Tmoxide. — 
Nitrogen  Dioxide  and  Nitrogen  Tetroxide.  —  Nitric  Oxide. 
—  Nitrous  Oxide.  —  Hyponitrous  Acid.  —  Exercises. 


CONTENTS.  X1H 

CHAPTER  XIII. 

PAGE. 

Sulphur  and  its  Compounds .105 

Occurrence  and  Preparation  of  Sulphur.  —  Physical  Proper- 
ties. —  Chemical  Properties.  —  Uses.  —  Compounds.  —  Hy- 
drogen Sulphide.  —  Properties.  —  Sulphides.  —  Precipitation 
of  Sulphides.  —  Carbon  Bisulphide.  —  Manufacture  of  Sul- 
phuric Acid.  —  Its  Purification.  —  Properties.  —  Reduction.  — 
Uses. —Sulphates. —Sulphur  Trioxide.  —  Sulphur  Dioxide. 

—  Chemical    Properties    of    Sulphur    Dioxide.- — Sulphurous 
Acid.  —  Thiosulphates.  —  Exercises. 

CHAPTER  XIV. 

Carbon  and  its  Compounds 186 

Carbon.  —  The  Diamond.  —  Graphite.  —  Amorphous  Car- 
bon.—  Coal. — Artificial  Amorphous  Carbon.  —  Carbon  Di- 
oxide; Occurrence.  —  Preparation.  —  Physical  Properties.  — 
Chemical  Properties.  —  Other  Sources  and  Uses.  —  Its  Rela- 
tion to  Life.  —  Carbonic  Acid.  —  Carbonates.  —  Bicarbonates. 
—  Natural  Carbonates.  —  Formation  of  Carbon  Monoxide.  — 
Laboratory  Method.  —  Properties.  —  Cyanogen.  —  Hydrocy- 
anic Acid.  —  Compounds  of  Carbon  and  Hydrogen.  —  Meth- 
ane. —  Ethane.  —  Ethylene.  —  Acetylene.  —  Illuminating  Gas. 

—  Illuminating    Gas   by    Distillation    of    Coal.  —  Water-Gas 
-Process.  —  Comparison     of     the     Two     Kinds     of    Gas.  — 
Amount  Used.  —  Exercises. 


CHAPTER   XV. 

Flames.     Heat  of  Formation  and  Decomposition 214 

Flames.  —  Luminosity.  —  Structure.  —  Non-Luminous 
Flames. —  Simple  and  Complex  Flames. — Dissection  of  Flames. 
—  Energy  Changes  Accompany  Chemical  Changes.  —  Heat  of 
Formation  and  Decomposition.  —  Positive  and  Negative  Heat 
of  Formation.  —  Heat  of  Formation  Evolved  in  Stages. 


XIV  CONTENTS. 

CHAPTER  XVI. 

PAGE. 

of  Multiple  Proportions;  the  Atomic  Theory;  Molecular 
and  Atomic  Masses,  etc 223 

Law  of  Multiple  Proportions.  —  The  Atomic  Hypothesis.  — 
Law  of  Definite  Proportions  Explained  by  the  Atomic 
Theory.  —  Explanation  of  Law  of  Multiple  Proportions.  — 
Distinction  between  Atoms  and  Molecules.  —  Molecular  Masses 
of  Gaseous  Substances.  —  Vapor  Density  Methods.  —  Other 
Methods.  —  Boiling  Point  and  Freezing  Point  Methods.  — 
Methods  of  Obtaining  Exact  Molecular  Masses.  —  Atomic 
Masses.  —  Their  Determination.  —  Exact  Atomic  Masses.  — 
Application  of  Atomic  Mass  Methods.  —  How  Formulas  are 
Determined.  — Molecular  Formulas  and  Equations.  — Nascent 
or  Atomic  State.  —  Number  of  Atoms  in  the  Molecules  of 
Elements.  —  Reason  for  Choosing  32  as  Molecular  Mass  of 
Oxygen.  —  Laws  of  Simple  and  Multiple  Volumes.  —  Volu- 
metric Meaning  of  an  Equation.  — Valence.  —  Different  For- 
mula Types  Based  on  Valence.  —  Graphic  Formulas.  — 
Isomerism.  —  Allotropism.  —  Exercises. 


CHAPTER  XVII. 


/Fluorine,  Bromine,  Iodine,  and  their  Compounds 


Acid. 


Halogens.  —  Fluorine.  —  Properties.  —  Hydrofluoric, 

—  Preparation     of     Bromine.  —  Properties.  —  llydrobromk' 
Acid.  —  Properties.  —  Occurrence  and  Preparation  ^HJfodine. 

—  Properties.  —  Hydriodic  Acid.  —  Properties.  —  Compounds 
of  the  Halogens  with  Oxygen.  —  With  Oxygen  aiyrllydrogen. 

—  Hypochlorous  Acid.  —  Chlorous    Acid.  —  Chloric   Acid.  — 
Perchloric  Acid.  —  Compounds  of  Bromine  with  Oxygen  and 
Hydrogen.  —  Compounds  of  Iodine  with  Oxygen  and  Hydro- 
gen. —  The  Halogen  Family.  — Exercises. 


CONTENTS.  XV 

CHAPTER  XVIII. 

PAGE. 

Ozone  and  Hydrogen  Peroxide 272 

Ozone.  —  Properties.  —  Hydrogen   Peroxide.  —  Properties. 
—  Composition.  —  Peroxides  and  Dioxides. 


CHAPTER  XIX. 

The  Nitrogen  Family 278 

PHOSPHORUS.  —  Occurrence  and  Preparation.  —  Properties. 
—  Red  Phosphorus.  —  Molecular  Mass  of  Phosphorus.  — 
Matches.  —  Hydrogen  Phosphide.  —  Phosphonium  Salts.  — 
Phosphides.  —  Oxides  of  Phosphorus.  —  Hypophosphorous 
Acid.  —  Phosphorous  Acid.  —  Phosphoric  Acids.  —  Prepara- 
tion. —  Salts.  — Uses  of  the  Phosphates. 

ARSENIC.  —  Occurrence  and  Preparation.  —  Properties.  — 
Hydrogen  Arsenide.  —  Marsh's  Test. — Arsenic  Trioxide. — 
Arsenious  Acid.  —  Arsenic  Acid.  —  Arsenic  Trisulphide. 

ANTIMONY.  — Preparation.  —  Properties.  —  Antimony  Com- 
pounds. —  Uses. 

BISMUTH.  —  Preparation.  —  Properties.  —  Bismuth  Salts.  — 
Uses.  —  The  Nitrogen  Family.  —  Exercises.  —  Table. 


CHAPTER  XX. 

The  Periodic  System 305 

Natural  Families.  —  Periodic  Arrangement.  —  Properties  of 
.  an    Element    Determined.  —  Gaps    in    the    Arrangement.  — 
Prediction  of  Unknown   Elements.  —  Conclusion.  —  Periodic 
Table.  —  Argon  Family. 

CHAPTER  XXI. 

Silicon  and  Boron 315 

Occurrence  of  Silicon. — Preparation. — Silicon  Compounds. 
—  Silicon     Dioxide.  —  Silicic     Acid.  —  Silicates.  —  Glass.  — 


X 


• 


CONTENTS. 

PAGE. 

Boron;  its  Occurrence,  Preparation,  and  Properties.  —  Boric 
Acid.  —  Borax.  —  Exercises. 


CHAPTER  XXTL 

Dissociation  and  Mass  Action 325 

Dissociation  by  Heat. — lonization. — Explanation  of  Neu^ 
tralization.  —  Test  Reactions  are  Ionic.  —  Electrolysis.  — 
Hydrolysis.  —  Mass  Action.  —  Exercises. 


CHAPTER  XXIII. 
334 

Metals  and  Non-Metals.  —  Occurrence  of  Metals.  —  Extrac- 
tion and  Properties  of  Metals. 

CHAPTER  XXIV. 

he  Alkali  Metals .     .337 

General  Properties.  —  Lithium. — Sodium;  its  Preparation 
and  Properties.  —  Sodium  Hydroxide.  —  Soap.  —  Sodium  Car- 
bonate; Bicarbonate;  Phosphate;  Chloride,  and  Nitrate. — 
Other  Salts.  —  Potassium. — Potassium  Hydroxide;  Carbon- 
ate; Nitrate;  Chlorate;  Bromide  and  Iodide. — Ammonium. 
—  Exercises. 


CHAPTER  XXV. 

The  Alkaline-Earth  Metals 350 

The  Group.  —  Calcium.  —  Calcium  Oxide;  Hydroxide; 
Chloride;  Carbonate;  Sulphate,  and  Phosphate.  —  Fertilizers. 
—  Mortar  and  Cement.  —  Other  Calcium  Compounds. — 
Barium.  —  Strontium.  —  Magnesium.  —  Magnesium  Com- 
pounds. —  Glucinum.  —  The  Group  a  Natural  Family.  — 
Table.  —  Exercises. 


CONTENTS.  XVU 

CHAPTER   XXVI. 

PAGE. 
Zinc,  Cadmium,  and  Mercury 361 

The  Zinc  Group. — Zinc;  its  Metallurgy,  Properties,  and 
Uses. — Zinc  Compounds. — Cadmium. — Mercury;  its  Proper- 
ties and  Uses.  —  Mercury  Compounds. 

CHAPTER  XXVII. 

Copper,  Silver,  and  Gold 366 

Relation  of  Copper,  etc.,  to  the  Alkali  Metals.  —  Copper; 
Occurrence  and  Metallurgy.  —  Properties.  — Uses.  —  Copper 
Compounds.  —  Copper-plating.  —  Silver.  —  Extraction  of 
Silver  from  its  Ores. — Reverberatory  Furnace. — Properties 
of  Silver.  —  Uses. — Compounds  of  Silver.  —  Photography. 
—  Gold;  its  Occurrence  and  Metallurgy.  —  Purification. — 
Properties  and  Uses. 

'CHAPTER  XXVIII. 

^-Aluminum 378 

Occurrence  of  Aluminum.  —  Preparation.  —  Hall  Process. 
—  Properties  of  Aluminum.  —  Uses.  —  Aluminum  Com- 
pounds.—  Alums. —Porcelain  and  Stoneware. 


CHAPTER  XXIX. 

Iron,  Nickel,  and  Cobalt 384 

Occurrence  of  Iron.  —  Metallurgy.  —  Blast  Furnace.  —  Com- 
mercial Iron. —  Tempering. —  Manufacture  of  Steel.  —  Crucible 
Process.  —  Bessemer  Process.  —  Converter.  —  Open  Hearth 
Process.  —  Properties  of  Iron.  —  Iron  Oxides  and  Hydroxides. 
—  Iron  Sulphides,  Chlorides,  and  Sulphates.  —  Ferrocyanides 
and  Ferricyanides.  —  Nickel.  —  Cobalt. 


XV111  CONTENTS. 

CHAPTER    XXX. 

PAGE. 
•^ianganese  and  Chromium 392 

Manganese;  its  Occurrence  and  Properties. — Manganous 
Salts.  —  Oxides  and  Hydroxides  of  Manganese.  —  Potassium 
Permanganate.  —  Oxidation  by  Potassium  Permanganate. 

Chromium;  its  Occurrence  and  Preparation.  —  Oxides  and 
Hydroxides.  —  Chromous  and  Chromic  Salts.  —  Double  Nature 
of  Chromium.  —  Chromates  and  Bichromates.  —  Oxidation 
by  Chromates  and  Dichromates. 


CHAPTER  XXXI. 
Lead,  Tin,  and  Platinum 400 

Occurrence  and  Metallurgy  of  Lead.  —  Properties  and  Uses. 
—  Lead  Compounds.  —  Occurrence  and  Metallurgy  of  Tin.  — 
Properties  and  Uses.  —  Tin  Compounds.  —  Occurrence  of 
Platinum. — Its  Preparation  and  Properties.  —  Chlorplatinic 
Acid. 


CONTENTS. 


PART   II. 


LABORATORY  EXERCISES. 

EXPERIMENT.  PAGB. 

I.     The  Bunsen  Burner 1 

II.     Cutting  and  Bending  Glass  Tubing 3 

III.  Effect  of  Heat  upon  "  Red  Precipitate  "      ...  4 

IV.  Solution,  Filtration,  and  Evaporation      ....  5 
V.     Hydrogen     . 7 

VI.     Equivalent  of  Magnesium 10 

VII.  Oxygen 11 

VIII.  Kindling  Temperature *  .  13 

IX.  Action  of  Sodium  upon  Water 14 

X.  Water  of  Crystallization 15 

XI.  Efflorescence 16 

XII.  Deliquescence 16 

XIII.  Effect  of  Temperature  on  Solution.     Crystalliza- 

tion       17 

XIV.  Precipitation 18 

XV.     Constant  Proportions 18 

XVI.     Chlorine 20 

XVII.     Hydrochloric  Acid 22 

XVIII.     Properties  of  Acids 24 

XIX.     Properties  of  Bases 26 

XX.     Properties  of  Salts 27 

XXL     Neutralization 28 

XXII.     Normal  and  Acid  Salts 29 

XXIII.  Normal  and  Acid  Salts  Continued 31 

XXIV.  Nitrogen 32 

XXV.    Ammonia     .    .    .    . 33 

Si* 


XX 


CONTENTS. 


EXPERIMENT.  PAGE. 

XXVI.  Nitric  Acid 36 

XXVII.  Nitrites 37 

XXVIII.  Nitrogen  Tetroxide 38 

XXIX.  Nitric  Oxide 39 

XXX.  Nitrous  Oxide .     .  41 

XXXI.  Physical  Properties  of  Sulphur 42 

XXXII.  Chemical  Properties  of  Sulphur 43 

XXXIII.  Hydrogen  Sulphide 44 

XXXIV.  Sulphur  Dioxide 40 

XXXV.  Sulphuric  Acid 48 

XXXVI.  Sulphates^ .     .     .    >    .     .  49 

XXXVII.  Carbon 50 

XXXVIII.  Carbon  Dioxide,  I 51 

XXXIX.  Carbon  Dioxide,  II 52 

XL.  Reduction.     Effect  of  Heat  on  Carbonates ...  53 

XLI.  Flames 54 

XLII.  Weight  of  a  Liter  of  Oxygen 55 

XLIII.  Bromine 57 

XLIV.  Iodine  and  Hydriodic  Acid 58 

XLV.  Comparison  of  the  Halogen  Acids 60 

XL VI.  Hydrogen  Peroxide 61 

XLVII.  Phosphorus  and  Phosphoric  Acid 62 

XLVIII.  Arsenic 63 

XLIX.  Antimony 65 

L.  Bismuth 66 

Li.  Borax  and  Boric  Acid 66, 

LII.  lonization 68 

LIII.  Hydrolysis  and  Mass  Action 68 

LIV.  Sodium  Compounds 69 

LV.  Potassium  Compounds 71 

LVI.  Solubility  of  Potassium  Chloride 72 

LVII.  Ammonium  Amalgam.    Distinctions  between  the 

Alkali  Metals 73 

LVIII.  Calcium 74 

LIX.  Water  of  Crystallization  in  Gypsum 75 

LX,  Strontium  and  Barium  ....  ,76 


CONTENTS.  XXI 

EXPERIMENT.  PAGE. 

LXI.  Water  of  Crystallization  in  Barium  Chloride    .     .  77 

LXII.     Magnesium 78 

LXIII.     Zinc 79 

LXIV.     Equivalent  of  Zinc 81 

LXV.     Cadmium 81 

LXVI.     Mercury 82 

LXVII.     Copper 84 

LXVIII.     Silver 85 

LX1X.     Aluminum 86 

LXX.     Iron 87 

LXXI.     Nickel  and  Cobalt 90 

LXXII.     Manganese  Compounds 90 

LXXIII.     Chromium  Compounds 91 

LXXIV.     Lead 93 

LXJLV.  Tin                                                .......  94 


CHEMISTRY. 


INTRODUCTION. 

1.  Definition   of  Chemistry. — Chemistry   may   be 
defined  as  the  science  which  deals  with  the  different  kinds 
of  matter  and  their  transformations.     The  kinds  of  matter 
are  called    substances.     Many  substances  occur   ready- 
formed  in  the  earth ;  but  many  others  are  not  so  found, 
and   must  be  made  from  existing  bodies.     Chemistry, 
therefore,  consists  largely  of  a  description  of  the  ways  in 
which  substances  occur  in  nature,  of   the  methods  by 
which  they  may  be  produced  in  the  laboratory,  and  of 
the  properties,  or   characteristics,  by  which    they  may 
be  distinguished  from  one  another. 

Descriptive  Chemistry  alone,  however,  cannot  give  a 
connected  and  an  intelligent  view  of  the  whole  science ; 
this  can  result  only  from  a  study  both  of  the  laws  gov- 
erning chemical  phenomena,  and  of  the  most  important 
theories  that  men  have  suggested  to  explain  the  laws. 

2.  Importance  of  Chemistry.  --The  ideas  of  modern 
Chemistry  are  of  great  importance,  both  for  their  own 


2  INTRODUCTION. 

sake,  and  because  they  add  so  much  to  other  depart- 
ments of  knowledge ;  no  apology  is  therefore  needed  for 
the  presence  of  Chemistry  in  a  course  of  study.  We 
may,  however,  summarize  in  brief  form  the  reasons  why 
every  student  in  a  secondary  school  should  master  at 
least  the  elements  of  chemical  science :  — 


1.  A  good  course  of  laboratory  work  in  Chemistry  disciplines 
the  mind,  as    few   courses   can,   in   independent   and   honest 
observation  of  phenomena. 

2.  A  knowledge  of  this  science  is  necessary  for  the  intelli- 
gent study  of  other  natural  sciences,  such  as  geology,  astronomy, 
biology,  physiology,  etc.,  which  make  special  application  of  its 
general  ideas. 

3.  Chemistry  is  intensely  practical,  for  its  facts  are  in   the 
most  common  use  in  the  arts  and  in  e very-day  life. 

As  illustrations  of  the  practical  character  of  chemical  knowl- 
edge we  may  cite  its  application  in  medicine,  in  sanitation,  in 
domestic  science,  in  the  extraction  of  metals  from  their  ores, 
in  the  refining  of  petroleum,  and  in  the  manufacture  of  steel, 
illuminating  and  fuel  gas,  paints,  dyestuffs,  food  products,  ice, 
alcohol,  soap,  glass,  paper,  explosives,  etc. 

To  be  sure,  it  is  impossible  for  an  elementary  course  in 
Chemistry  to  give  all  the  facts  relating  to  such  topics  as  those 
suggested  above  ;  but  it  does  give  the  facts  and  the  theories 
that  are  fundamental,  and  by  means  of  which  even  more  com- 
plex phenomena  must  be  interpreted. 


3.  Relation  between  Chemistry  and   Physics.  - 

Chemistry  and  Physics  are  closely  related  sciences,  for 
both  together  have  as  their  object  the  study  and  the 

\       •         "'  '      .Vli-'IU^ 

I  ; & •  ':- 


INTRODUCTION.  3 

explanation  of  the  general  phenomena,  or  changes,  that 
take  place  in  the  material  universe.  They  are,  in  fact, 
two  points  of  view  from  which  natural  phenomena 
may  be  considered.  Physics  has  to  do  chiefly  with 
transformations  of  energy,  and  with  matter  only  as 
that  upon  which  energy  acts  to  produce  phenomena. 
Uhemistry,  on  the  other  hand,  is  largely  concerned  with 
the  various  forms  of  matter,  and  with  energy-changes 
only  as  they  result  in  the  formation  of  new  sub- 
stances. 

Accordingly,  phenomena  are  usually  distinguished  as 
either  physical  or  chemical  phenomena.  To  the  former 
class  belong  those  changes  in  which  the  substance,  in 
connection  with  which  the  energy  manifests  itself,  is  not 
permanently  altered,  but  regains  its  original  properties. 
A  physical  change  may,  therefore,  be  repeated  with  the 
same  substance  after  the  substance  has  resumed  its 
former  condition.  Illustrations  are :  the  magnetization 
of  a  knife-blade;  the  production  of  light  by  means  of 
white-hot  iron ;  the  melting  of  ice,  and  the  vaporization 
of  water. 

Chemical  phenomena,  on  the  other  hand,  involve  a 
permanent  alteration  of  the  properties  of  the  substances 
used.  Thus,  a  piece  of  burnt  magnesium  does  not 
assume  its  original  condition  on  cooling ;  rusted  iron  is 
no  longer  iron ;  carbonic  acid  gas  is  neither  carbon  nor 
oxygen,  although  these  two  substances  were  put  together 
to  produce  it. 

In  the  same  way  we  distinguish  between  the  properties 


INTRODUCTION. 


of  substances,  calling  those  properties  physical  which 
require  only  physical  phenomena  for  their  exhibition, 
and  those  properties  chemical  which  are  capable  of  being 
manifested  only  by  some  chemical  change. 


Thus  the  color,  specific  gravity,  melting  point,  crystalline 
form,  solubility,  etc.,  of  sulphur  would  be  termed  physical 
properties  of  sulphur  ;  its  power  of  burning  in  air  is,  on  the 
contrary,  a  chemical  property. 


The  description  of  a  substance  in  Chemistry  includes 
its  most  important  physical,  as  well  as  its  chemical, 
properties. 

4.  Reagents  and  Reactions.  —  Substances  which 
have  a  chemical  effect  upon  one  another  are  said  to 
react,  and  a  chemical  change  is  therefore  called  a  re- 
action.  The  substances  which  react  are  called  reagents, 
or  factors.  Thus,  when  copper  is  treated  with  concen- 
trated nitric  acid,  a  chemical  reaction  takes  place,  and 
the  copper  and  the  nitric  acid  are  the  reagents  (or 
factors). 

The  adjective  "  reagent "  is  often  applied  to  sub- 
stances in  the  form  in  which  they  are  commonly  used  in 
the  laboratory.  Thus  "  reagent "  ammonia  means  an 
aqueous  solution  of  ammonia ;  ammonia  itself  is  a  gas. 
Similarly,  by  "  reagent  "  sodium  hydroxide  we  mean  the 
solution  of  solid  sodium  hydroxide  in  water,  this  being 


INTRODUCTION.  5 

the  form  in  which  sodium  hydroxide  is  most  frequently 
used  in  the  laboratory. 

5.  Elements  and  Compounds. — Almost  all  01  tne 
substances  found  in  the  earth  may  be  shown,  by  one 
method  or  another,  to  consist  of  two  or  more  different 
kinds  of  matter,  and  are  therefore  called  compound 
substances,  or  compounds.  There  are,  however,  between 
seventy  and  eighty  substances  which  it  is  impossible  for 
us  to  decompose,  with  our  present  methods,  into  simpler 
kinds  of  matter ;  these  substances  are,  therefore,  called 
elementary  substances,  or  simply  elements. 

Less  than  half  of  the  substances  believed  to  be  elementary 
are  really  found  free  in  the  earth  and  its  atmosphere  ;  the 
others  occur  only  in  a  combined  form. 

Although  the  number  of  elements  is  so  small,  the 
number  of  compounds  they  are  actually  known  to  form 
is  very  large  —  probably  not  less  than  one  hundred 
thousand.  The  number  of  compounds  theoretically  pos- 
sible, but  not  yet  known  to  exist,  is  also  very  large. 

A  list  of  the  substances  usually  considered  elementary  follows. 
The  letter,  or  combination  of  letters,  given  after  the  name  of 
each  element  is  called  the  symbol  of  the  element.  It  is  not  in- 
tended that  all  these  symbols  shall  be  learned  now,  but  rather 
by  association  with  the  names  of  the  elements  as  the  latter  are 
studied.  In  cases  in  which  symbols  are  formed  from  the  Latin 
(or  Greek)  names  of  the  elements  instead  of  the  English  names, 
both  names  are  given. 


6 


INTRODUCTION. 


ELEMENTS. 

SYMBOLS. 

ELEMENTS. 

SYMBOLS. 

Aluminum. 

Al. 

Neodymium. 

Nd. 

Antimony. 

Sb. 

Nickel. 

Ni. 

Argon. 

A. 

Niobium. 

Nb. 

Arsenic. 

As. 

Nitrogen. 

N. 

Barium. 

Ba. 

Osmium. 

Os. 

Beryllium  (Glucinum). 

Be.  (Gl.) 

Oxygen. 

O. 

Bismuth. 

Bi. 

Palladium. 

Pd. 

Boron. 

B. 

Phosphorus. 

P. 

Bromine. 

Br. 

Platinum. 

Pt. 

Cadmium. 

Cd. 

Potassium  (Kallum). 

K. 

Caesium. 

Cs. 

Praseodymium. 

Pr. 

Calcium. 

Ca. 

Rhodium. 

Rh. 

Carbon. 

C. 

Rubidium. 

Rb. 

Cerium. 

Ce. 

Ruthenium. 

Ru. 

Chlorine. 

Cl. 

Samarium. 

Sm. 

Chromium. 

Cr. 

Scandium. 

Sc. 

Cobalt. 

Co. 

Selenium. 

Se. 

Copper  (Cuprum). 

Cu. 

Silicon. 

Si. 

Erbium. 

Er. 

Silver  (Argeutum). 

Ag. 

Fluorine. 

Fl. 

Sodium  (Natrium). 

Na. 

Gallium. 

Ga. 

Strontium. 

Sr. 

Germanium. 

Ge. 

Sulphur. 

S. 

Gold  (Aurum). 

Au. 

Tantalum. 

Ta. 

Helium. 

He. 

Tellurium. 

Te. 

Hydrogen. 

H. 

Thallium. 

Tl. 

Indium. 

In. 

Thorium. 

Th. 

Iodine. 

I. 

Tin  (Stannum). 

Sn. 

Iridium. 

Ir. 

Titanium. 

Ti. 

Iron  (Ferrum). 

Fe. 

Tungsten  (Wolframium). 

W. 

Lanthanum. 

La. 

Uranium. 

Ur.  ' 

Lead  (Plumbum). 

Pb. 

Vanadium. 

Vd. 

Lithium. 

LI. 

Ytterbium. 

Yb. 

Magnesium. 

Mg. 

Yttrium. 

Y. 

Manganese. 

Mn. 

Zinc. 

Zn. 

Mercury  (Hydrargyrum). 

Hg. 

Zirconium. 

Zr. 

Molybdenum. 

Mo. 

Besides  the  elements  given  in  the  above  list,  there  are 
several  other  substances  which  are  considered  by  some 


INTRODUCTION.  7 

chemists  to  be  elementary,  but,  the  true  nature  of  these 
substances  being  still  in  dispute,  their  names  are 
omitted. 

6.  Relative  Abundance  and  Importance  of  the 
Elements.  —  The  elements  are  by  no  means  equally 
abundant ;  nor  are  they  all  of  equal  importance  for  the 
organic  forms  existing  upon  the  earth.  It  is  probable 
that  only  eleven  are  absolutely  essential  to  animal  life. 
These  are,  — 

Carbon,  Sulphur,  Phosphorus, 

Oxygen,  Calcium,  Potassium, 

Nitrogen,  Sodium,  Iron. 

Hydrogen,  Chlorine, 

If  four  more  were  present,  viz. :  — 

Silicon,  Magnesium, 

Aluminum,  Fluorine, 

savage  life,  upon  an  earth  similar  to  ours,  would  be 
possible. 

With  seven  additional  elements,  viz. :  — 

Gold,  Tin,  Manganese, 

Silver,  Zinc,  Mercury, 

Platinum, 

modern  civilization  might  exist. 

The  great  inequality  in  the  distribution  of  the  ele- 
ments is  shown,  in  a  rough  way,  by  Fig.  1.  As  there 


8 


INTRODUCTION. 


indicated,  silicon  and  oxygen  together  make  up  about 
three  fourths  of  the  earth's  solid  crust. 


ttf* 

Potassium  l./°      . 
All  others  not    \./<> 


FIG.  i. 

We  shall  begin  our  study  of  Chemistry  in  Chapter  I 
with  the  element  that  is  in  some  respects  the  most 
remarkable  of  all,  —  the  element  hydrogen. 

7.  Exercises. 

Classify  the  following  as  either  physical  or  chemical  phe- 
nomena :  — 

The  souring  of  milk,  the  burning  of  wood,  the  evaporating  of 
water,  the  tarnishing  of  silver,  the  dissolving  of  sugar  in  water, 
the  bleaching  of  muslin,  the  melting  of  lead. 


CHAPTER   I. 
HYDROGEN. 

8.  Existence.  —  Hydrogen  is  a   light,  colorless   gas 
that   is    found    in   a  free  condition    only   in   compara- 
tively small  amounts  on  the  earth  —  chiefly  in  the  air 
and  in  volcanic  gases.     It   exists  in  great   quantities, 
however,  in  the  atmosphere  of  the  sun.     Although  rare 
in  the  uncombined  form,  hydrogen  is  a  constituent  of 
many   important   and    abundant   compounds,    such    as 
organic  substances,  water,  acids,  etc.     The  most  com- 
mon compound  of  hydrogen  is  water.     One  ninth,  by 
weight,  of  water   is   hydrogen,    and    the  remainder   is 
another  colorless  gas,  viz.,  oxygen. 

The  name  "hydrogen"  means  "  a  producer  of  water"; 
while  "  oxygen  "  means  "  a  producer  of  acids." 

9.  Common  Method  of  Preparation.  —  Since  water 
is  a  compound  of  hydrogen,  the  decomposition  of  water 
will,  of  course,  give  hydrogen ;  but  a  much  more  con- 
venient way  to  prepare  the  gas  is  to  decompose  certain 
acids.     Acids,  like  water,  contain  hydrogen,  and  give  it 
up  readily  when  treated  with  certain  metals  under  ap- 
propriate   conditions.      The    metal    commonly    used   is 
zinc,  and   the   acid    either   hydrochloric  or  dilute  sul- 
phuric acid, 


10 


HYDROGEN. 


Generation  and  Collection  of  Hydrogen. 

The  gas  is  usually  produced  in  a  generating  flask  (Fig.  2) 
provided  with  a  "  thistle  "  or  "  safety  "  tube  and  a  delivery 
tube  reaching  into  the  water  of  a  water  pan  (called  also  a 
"  pneumatic  trough ").  The  flask  contains  zinc.  Acid  is 
added  through  the  thistle  tube  until  the  lower  end  of  the 
thistle  tube  is  immersed  ;  the  evolved  gas  escapes  through  the 


FIG.  2. 

delivery  tube  and  bubbles  up  through  the  water  in  the  pan. 
Here  the  hydrogen  may  be  collected  in  appropriate  "  receiv- 
ers "  (test  tubes,  bottles,  etc.)  filled  with  water  and  inverted 
over  the  end  of  the  delivery  tube.  Or,  since  hydrogen  is  much 
lighter  than  air,  it  may  be  collected  by  displacing  the  air  of  the 
receiver,  as  is  shown  in  Fig.  3. 

N.  B.  Apparatus  in  which  hydrogen  is  being  gener- 
ated must  not  be  brought  near  a  flame  ! 

If  the  action  between  metal  and  acid  is  not  brisk,  it 


PURIFICATION   OF  HYDROGEN. 


11 


may  be    hastened    by  adding   a    few   drops    of   copper 
sulphate  solution.     The  copper  sulphate  reacts  with  a 
portion  of  the  zinc,  precipitating  copper,  which  forms 
a   black    deposit    upon    the    zinc;    thus 
coated,  the  latter  acts  readily  upon  the 
acid.     The  action  of  zinc  and  acid  re- 
sults in   a  very  considerable    evolution 

of  heat. 

Self-Regulating  Generator. 

Instead  of  the  ordinary  generating  flask, 
a  Kipp's  or  other  self-regulating  apparatus 
may  be  used  to  supply  hydrogen.  Kipp's 
apparatus  (Fig.  4)  consists  of  three  globes. 
The  upper  globe  is  in  communication  with 
the  lower  globe,  and  the  middle  globe  with 
the  lower  globe,  but  the  upper  globe  and  the 
middle  globe  are  not  connected.  The 
upper  and  lower  globes  contain  dilute  acid, 
but  the  middle  globe  contains  zinc.  This 
is  the  condition  of  the  apparatus  when  at  rest,  with  the  stop- 
cock closed.  When  the  stopcock  is  opened,  the  liquid  of 
the  upper  globe  falls  into  the  lower  globe,  and  the  liquid  in 
the  lower  globe  rises  into  the  middle  globe,  thus  displacing 
the  gas  of  the  middle  globe,  and  forcing  it  out  through  the 
stopcock.  The  acid  which  enters  the  middle  globe  reacts 
with  the  zinc,  forming  more  hydrogen,  which  either  escapes 
through  the  stopcock,  or,  if  the  latter  is  closed,  forces  the  acid 
back  into  the  lower  globe  and  thence  into  the  upper  globe. 
The  gas  in  the  middle  globe  is  thus  ready  for  instant  use. 


FIG.  3. 


10.  Purification  of  Hydrogen.  —  If  the  metal  or  the 
acid  used  in  preparing  hydrogen  is  of    "  commercial " 


12 


HYDROGEN, 


grade,  the  hydrogen  will  be  impure,  as  may  be  inferred 
from  its  disagreeable  odor.     We  may  remove  most  of 
the  impurities  —  as  well  as  volatile  acid,  if  hydrochloric 
acid  is  used  —  by  passing  the  gas 
through  a  mixture  of   sodium  hy- 
droxide  and  potassium  permanga- 
nate solutions.     To  dry  it  we  use 
some  dehydrating  agent,  e.  g.,  gran- 
ular calcium  chloride. 

The  apparatus  for  preparing 
comparatively  pure  hydrogen  is 
shown  in  Fig.  5. 


The  hydrogen  is  generated  in  a  flask 
(or  Kipp's  apparatus),  and  is  passed 
through  a  bottle  containing  potassium 
permanganate  dissolved  in  10%  sodium 
hydroxide  solution,  and  then  through 
a  U-tube  of  calcium  chloride.  A  sec- 
ond bottle  of  the  permanganate  solu- 
tion will  make  the  purification  more 
complete.  The  exit  tube  of  the  U- 
tube  is  drawn  out  to  a  small  opening, 
so  that  the  hydrogen  shall  issue  in  a 
steady  stream.  If  the  washing  of  the 
hydrogen  by  the  "  alkaline  perman- 
ganate "  solution  has  been  success- 
ful, the  gas  will  now  be  practically 
odorless. 


FIG.  4. 


The  stream  of  gas  may  be  lighted  if  the  following 
precautions  are  observed :  — 


PURIFICATION   OF  HYDtiOGEN. 


13 


In  every  case,  before  we  light  a  jet  of  hydrogen  (the 
same  precautions  apply  to  other  inflammable  gases),  we 
collect  a  test  tube  full  by  displacing  the  air,  and  then  carry 
the  test  tube  —  in  this  case,  mouth  down  —  to  a  gas  jet 
or  other  flame  at  least 
four  feet  away.  The 
gas  in  the  test  tube 
is  thus  set  on  fire, 
with  explosion,  if 
there  is  still  much  air 
mixed  with  the  hy- 
drogen, but  quietly, 
if  the  air  originally 
in  the  apparatus 
has  been  displaced. 
We  then  carry  the 
test  tube  back  to  FIG.  5. 

the  stream    of   hy- 
drogen and  repeat  the  operation  with  freshly  collected 
test  tubes  full  of   the  gas,  until  the  hydrogen  stream 
is  lighted. 

The  reason  for  all  this  precaution  is  that  it  is  unsafe 
to  light  a  confined  mixture  of  hydrogen  and  air;  and  we 
can  be  sure  that  the  displacement  of  the  air  in  the 
apparatus  is  complete  only  when  the  time  needed  for  the 
burning  of  the  test  tube  of  hydrogen  is  greater  than 
that  required  for  the  return  of  the  test  tube  to  the 
jet  of  hydrogen. 

The  action  of  zinc  upon  dilute  sulphuric  acid  gives, 


14  HYDROGEN. 

besides  hydrogen,  a  substance  called  zinc  sulphate. 
This  remains  in  the  solution,  but  may  be  obtained  from 
it  as  a  white,  crystalline  solid. 

ii.  Chemical  Properties. — The  hydrogen  flame  is 
almost  colorless,  but  very  hot,  as  holding  a  piece  of 
platinum  in  it  will  show.  Indeed,  there  is  enough  heat 
liberated  by  the  burning  of  a  gram  of  hydrogen  in  oxy- 
gen gas  to  raise  the  temperature  of  about  34,000  grams 
of  water  1°  C.,  or  about  340  grams  from  the  freezing 
point  to  the  boiling  point. 

An  apparatus  for  making  use  of  this  great  heat  evolution  is 
the  oxyhydrogen  blowpipe  (Fig.  6).  This  consists  of  a  small 

inner  tube  communicating  with  a 
0  *"  *r?  ,-  ^  tank  of  oxygen,  and  a  larger  outer 

tube  in  connection  with  a  tank  of 

hydrogen.  Both  gases  are  greatly 
H  compressed.  The  hydrogen  is 

FIG.  6.  first  turned  on  and  lighted  ;  then 

the  oxygen  is  allowed  to  escape 

very  slowly.  Thus  a  flame  is  produced  which  is  so  hot  that  it 
will  melt  platinum.  (Platinum  melts  at  about  1700°  C.)  A 
piece  of  quicklime  held  in  the  oxyhydrogen  flame  becomes 
white  hot  and  gives  off  much  light ;  this  is  the  so-called  calcium, 
or  Zime,  light.  In  the  production  of  the  lime  light  for  stereopti- 
cons,  illuminating  gas  is  generally  used  instead  of  hydrogen. 
The  ordinary  blast-lamp  of  laboratories  is  similar  to  the  oxy- 
hydrogen blowpipe,  but  the  gases  used  are  illuminating  gas 
and  ordinary  air  ;  as  a  result,  the  temperature  produced  is  by 
no  means  as  high  as  that  of  the  oxyhydrogen  flame. 

The  hydrogen  flame  has  only  one  zone,  or  region,  of 


UNION  OF  HYDROGEN   WITH  OXYGEN.  15 

combustion.     A  vertical  section  of    it  would    have  the 
appearance  shown  in  Fig.  7,  - 

a  is  the  central  space  of  unhurried  hydrogen ; 

b  is  the  region  of  combustion. 

The  bearing  of  this  fact  will  be  understood 
when  Ave  come  to  study  more  complex  flames. 

12.  Union  of  Hydrogen  with   Oxygen.  - 

The  product  formed  when  hydrogen  burns  in  air       FIG>  7 
is  water,  as  we  may  prove  readily  by  holding 
over  the    burning    hydrogen   a    beaker    of    cold  water 
(Fig.  8).     Although  the  beaker  was  dry  on  the  outside 
at  the  beginning  of  the  experiment,  it  will  soon  con- 
dense drops  of  water  from  the  flame.     The  water  may 
be  easily   collected    and  identified. 

In  burning,  hydrogen  takes  oxygen  from 
the  air  ;  in  fact,  every  case  of  combustion  in 
air  consists  in  the  union  of  the  substance 
burned  with  the  oxygen  of  the  air.  Hydrogen 
does  not,  however,  support  combustion ;  that 
is  to  say,  burning  wood,  paper,  illumina- 
ting gas,  etc.,  —  the  ordinary  combustibles, 
—  do  not  continue  to  burn  when  placed  in 
an  atmosphere  of  hydrogen. 

FIG.  s.  Hydrogen  is  really  inert  at  ordinary  tem- 

peratures, and  active  only  when  the  temper- 
ature is  raised  considerably.  This  is  shown  by  the 
fact  that  even  hydrogen  and  oxygen  may  be  mixed 
and  kept  together  for  an  indefinite  period  without 


16  HYDROGEN. 

evidence  of  action.  The  inertness  of  hydrogen  under 
ordinary  conditions  is  such  that  it  is  often  conven- 
ient to  experiment  with  a  substance  in  an  atmosphere 
of  this  gas,  thus  excluding  the  active  oxygen  of  the  air. 

At  about  350°  C.,  however,  hydrogen  and  oxygen  unite 
with  great  violence.  Although  only  a  very  8mall  amount 
of  the  mixture  of  the  gases  is  really  heated  to  850°  C.  by 
the  match  or  electric  spark  used  to  start  the  combustion,  yet 
the  burning  of  this  portion  causes  so  much  heat  to  be  given 
off  that  adjacent  portions  are  quickly  raised  to  the  required 
temperature ;  as  a  result,  union  takes  place  very  rapidly 
through  the  whole  mixture.  The  explosion  seems  instanta- 
neous, but  is  not ;  its  rate  has  been  determined  by  photog- 
raphy to  be  a  little  less  than  two  miles  a  second. 

If  the  union  of  oxygen  and  hydrogen  takes  place  at  a 
pressure  lower  than  the  ordinary  atmospheric  pressure,  the 
rate  of  the  explosion  and,  consequently,  its  violence,  is  much 
diminished. 

Hydrogen  has  the  power  of  combining  not  only  with 
free  oxygen,  but  also,  in  many  cases,  with  oxygen  that 
is  in  combination  with  other  elements.  Thus,  if  hydro- 
gen is  passed  over  heated  copper  oxide  and  lead  oxide, 
it  unites  with  the  oxygen  of  these  substances,  forming 
water,  and  setting  free  copper  and  lead,  respectively. 
These  are  illustrations  of  the  reducing,  i.  e.,  deoxidizing, 
power  of  hydrogen. 

13.  Occlusion  of  Hydrogen.  —  A  remarkable  prop- 
erty of  hydrogen  is  its  absorption  in  large  quantities  by 
certain  metals,  e.  #.,  platinum  and  palladium.  The 


PHYSICAL   PROPERTIES.  17 

hydrogen  is  said  to  be  "  occluded "  by  these  metals. 
The  phenomenon  may  be  illustrated  by  holding  a  piece 
of  platinum  sponge  in  a  jet  of  dry  hydrogen ;  so  much 
heat  is  evolved  during  the  occlusion  of  the  hydrogen 
that  the  latter  is  set  on  fire. 

The  ready  union  of  hydrogen  and  other  inflammable  gases 
with  air,  in  the  presence  of  platinum,  may  be  shown  by  directing 
the  mixture  of  cold  gases  issuing  from  a  Bunsen  burner 
against  a  piece  of  hot  platinum  foil.  The  gases  will  unite 
with  so  great  an  evolution  of  heat  that  the  platinum  will  con- 
tinue to  glow. 

Palladium  has  an  even  greater  power  of  occluding 
hydrogen  than  platinum  has.  One  volume  of  palladium 
can  absorb  about  375  volumes  of  hydrogen  at  the 
ordinary  temperature  I 

14.  Physical  Properties. — As  we  have  already 
learned,  hydrogen  is  a  colorless  and  odorless  gas.  It  is 
the  lightest  substance  known,  air  being  14  and  oxy- 
gen 16  times  as  heavy.  The  weight  of  one  liter  of 
hydrogen  at  0°  C.  and  760  mm.  pressure  is  about  .09 
gram ;  there  are,  therefore,  no  less  than  11  liters  of  the 
gas  to  the  gram. 

Hydrogen  is  the  standard  of  density  in  the  case  of 
gases  ;  its  relative  density  is  1. 

The  rate  at  which  hydrogen  diffuses,  i.  e.,  mixes  with 
other  gases,  is  four  times  that  of  oxygen.  A  special 
form  of  diffusion,  viz.,  transpiration,  may  be  illustrated 
as  follows :  — 


18  HYDROGEN. 

A  porous  cup  (Fig.  9)  is  attached  securely  to  a 
glass  tube  ending  under  water.  If,  now,  a  bell-jar  or 
a  large  bottle  filled  with  hydrogen  is  placed  over, 
and  enclosing,  the  porous  cup,  bubbles  of  gas  will  be 
seen  escaping  from  the  lower  end  of  the  tube. 

The  explanation  of  the  phenomenon  is  that 
the  two  gaseous  media,  air  and  hydrogen,  sepa- 
rated by  the  porous  partition,  tend  to  form  a 
FIG.  9.  homogeneous  mixture.  But  the  rate  at  which 
the  hydrogen  passes  through  the  partition  is  so  much 
greater  than  that  of  the  air,  that  an  increase  of  volume, 
and,  therefore,  of  pressure,  occurs  within  the  porous 
cup.  Consequently  some  of  the  gaseous  mix- 
ture escapes.  When  the  bell-jar  is  removed, 
the  reverse  diffusion  takes  place. 

A  somewhat  simpler  form  of  apparatus  is  shown 
in  Fig.  10.  A  wide  glass  tube  has  one  end  cov- 
ered with  a  cap  of  plaster  of  Paris.  If  the  tube 
is  filled  with  hydrogen  by  displacing  the  air,  and  the 
open  end  is  placed  at  once  under  water,  water  will 


FIG.  10. 
rise  in  the  tube. 

Hydrogen  is  not  very  soluble  in  water ;  at  14°  C. 
100  c.c.  of  water  absorb  only  1.9  c.c.  of  the  gas. 

At  the  ordinary  temperature,  hydrogen  cannot  be 
liquefied  by  any  pressure,  however  great;  but  by  intense 
cold,  in  addition  to  great  pressure,  the  gas  has  been  con- 
densed to  the  liquid  and  solid  state.  Liquid  hydrogen 
is  a  colorless  substance  less  than  one  tenth  as  heavy  as 
water.  It  boils  at  about  — 240°  C.,  at  the  ordinary 
pressure. 


CHEMICAL    CHANGES.  19 

Hydrogen  is  a  better  conductor  of  heat  and  of  elec- 
tricity than  any  other  gas. 

15.  Other  Methods  of  Preparation.  —  The  action  of 
metals  upon  acids  is  only  one  of  many  methods  by  which 
hydrogen  may  be  prepared ;  other  general  methods  are 
the  following :  — 

(a)  The  decomposition  of  water  by  the  electric  current. 
This  operation  is  called  the  "  electrolysis  "  of  water ;  it  takes 
place    only  when   the   water    contains   small   quantities  of 
certain  substances  called  "  electrolytes.'1     The  electrolyte 
commonly  used  is  dilute  sulphuric  acid. 

(b)  The  action  of  certain  substances  —  chiefly  metals  — 
upon  water.     Some  metals,  e.  </.,  sodium  and  potassium,  act 
violently  even  upon  cold  water;  others,  like  magnesium, 
zinc,  and  iron,  decompose  only  hot  water,  or  steam. 

(c)  The  action   of   metals   upon  basic  hydroxides,  e.  g., 
aluminum  filings  upon  sodium  hydroxide  solution. 

16.  Quantitative  Character  of  Chemical  Changes. 
—  When  hydrogen  unites  with  oxygen  to  form  water,  the 
relation  between  the  masses  of  the  elements  entering  into 
combination  is  a  definite  and  constant  one. 

Eight  parts,  by  weight,  of  oxygen  combine  with  one  part  of 
hydrogen. 

The  relation  between  the  mass  of  hydrogen  burned 
and  that  of  the  water  formed  is  also  constant ;  one  part 
of  hydrogen  always  forms  by  its  combustion  nine  parts 
of  water. 


20  HYDROGEN. 

Conversely,  nine  parts  by  weight  of  water  give,  when  decom- 
posed, eight  parts  of  oxygen  and  one  part  of  hydrogen. 

Similarly,  an  exact  relation  exists  between  the  masses 
of  zinc  and  sulphuric  acid  which  react  with  each  other, 
and  between  each  of  these  and  the  masses  of  hydro- 
gen and  of  zinc  sulphate  formed.  This  relation  is  as 
follows :  — 

Sixty-five  parts  by  weight  (grams,  pounds,  etc.)  of  zinc  and 
98  parts  of  sulphuric  acid  (diluted  with  water)  give  161  parts 
of  zinc  sulphate  and  two  parts  of  hydrogen. 

It  may  be  proved  by  the  most  careful  weighing  that 
when  zinc  reacts  with  sulphuric  acid  there  is  no  gain  or 
loss  of  matter,  but  only  the  liberation  of  2  grams  of 
hydrogen  by  every  65  of  zinc.  If  the  quantities  of  zinc 
and  sulphuric  acid  taken  are  in  exactly  the  correct 
proportion,  and  the  resulting  solution  is  evaporated, 
neither  zinc  nor  sulphuric  acid  will  remain ;  and  the 
zinc  sulphate  obtained  will  contain  all  the  matter  of  the 
zinc  and  of  the  acid  used  except  the  hydrogen,  which 
escaped  from  the  solution. 

17.  Calculation  of  Quantities  of  Factors  and  Prod- 
ucts. —  Since  the  reaction  between  zinc  and  sulphuric 
acid  takes  place  in  definite  proportions  by  weight,  it  is 
always  possible  for  us  to  calculate  how  much  acid  will 
be  required  to  react  exactly  with  a  given  mass  of  zinc, 
and  how  much  hydrogen  and  zinc  sulphate  will  be 
formed. 


is- 

1* 

CALCULATION   OF   QUANTITIES.  21 

Suppose,  for  example,  that  we  wish  to  know  how  much 
sulphuric  acid  is  needed  to  react  exactly  with  40  grams  of 
zinc. 

Since  65  grams  of  zinc  require  98  grams  of  sulphuric  acid 
for  complete  action,  40  grams  of  zinc  will  require  a  propor- 
tional amount  of  the  acid.  The  amount  needed  may,  therefore, 
be  determined  b}'  solving  the  proportion,  — 

65  :  40  :  :  98  :  x 
whence  x  =  60.3  grams. 

This  is,  of  course,  the  least  quantity  of  sulphuric  acid 
that  will  use  up  all  of  the  zinc.  If  a  larger  amount  is 
used,  the  excess  will  simply  remain  in  the  solution. 

If  we  wish  to  find  the  amount  of  zinc  needed  to  pro- 
duce 5  grains  of  hydrogen,  we  may  do  so  by  solving  for 
x  in  the  proportion,  — 

65  :  x  :  :  2  :  5 
whence  x  —  162.5  grams. 

To  find  the  volume  at  0°  C.  and  760  mm.  pressure  of 
the  5  grams  of  hydrogen  produced,  we  divide  the  number 
of  grams  of  hydrogen  by  the  weight,  in  grams,  of  one 
liter  of  hydrogen  at  0°  C.  and  760  mm.  pressure  ;  i.  e.,  by 
0.09. 

The  volume  of  the  5  grams  of  hydrogen  is  thus,  — 


If  we  reverse  the  conditions  of  the  problem,  and  ask 


22  HYDROGEN. 

how  much  zinc  is  required  to  produce  by  its  action 
Avith  sulphuric  acid  40  liters  of  hydrogen  at  0°  C.  and 
760  mm.  pressure,  we  must  first  find  the  weight  of  the 
40  liters  of  hydrogen  at  0°  C.  and  760  mm.  pressure, 
and  then  calculate  the  weight  of  zinc  necessary. 

The  weight  of  40  liters  of  hydrogen  under  the  given  con- 
ditions is  3.6  (=  40  X  0.09)  grams  ;  the  weight  of  zinc  needed 
is  then  calculated  from  the  proportion,  — 

65  :  x  :  :  2  :  3.6. 

When  zinc  reacts  with  hydrochloric  acid  the  products 
are  zinc  chloride  and  hydrogen.  The  relative  quantities 
are  as  follows :  - 

65  grams  of  zinc  and  73  grams  of  hydrochloric  acid 
give  136  grams  of  zinc  chloride  and  2  grains  of  hydro- 
gen. 

1 8.  Exercises. 

1.  Learn  or  review  the  metric  tables  of  weight,  of  length, 
and  of  volume. 

2.  How  many  cubic  centimeters  in  1  liter?     In  1  c.  dm.? 
What  is  the  relation  between  g.,  mg.,  dg.,  eg.,  kg.  ?     What  is 
the   weight    of   106   c.c.    water?     State  the  relation  between 
1  mm.,  1  cm.,  1m.,  and  1  dm. 

3.  How    many  grams  of  hydrogen  can  be  stored  in  a  36- 
liter  gasometer  under  conditions  at  which  1  liter  of  hydrogen 
weighs  .09  grams? 

4.  What  properties  of  hydrogen  make  it  useful  as  the  in- 
flating gas  of  balloons  ?     What  properties  make  it  disadvan- 
tageous ? 

5.  How   many   grams   hydrogen   can    be   obtained    by   the 


EXERCISES.  23 

action   of   20  grams   zinc  upon  an  excess  of  dilute  sulphuric 
acid  ?     Upon  an  excess  of  hydrochloric  acid  ? 

6.  What  is  the  volume  in  cubic  centimeters  of  30  grams 
water?  Of  30  grams  platinum  of  S.G.  21.5?  Of  30  grams 
ether  of  S<G.  0.72? 


CHAPTER   II. 
OXYGEN. 

19.  Preparation.  — Oxygen  may  be  prepared  by  the 
electrolysis  of  water  (cf.  §  15),  but  for  ordinary  purposes 
the  decomposition  of  potassium  chlorate  is  more  con- 
venient. Potassium  chlorate  is  a  white,  crystalline 
solid  of  which  about  39^>  is  oxygen ;  all  of  this  oxygen 
is  given  off  upon  the  application  of  heat.  If  the  potas- 
sium chlorate  is  heated  by  itself,  a  test  tube  or  retort 


FIG.  11. 

(Fig,  11)  of  hard  glass  is  needed ;  for  the  temperature 
at  which  decomposition  begins  to  take  place  is  350°  to 
400°  C. 

There  are  two  stages  in  the  decomposition  of  the  potassium 
chlorate  When  it  is  first  heated,  potassium  chloride,  contain- 
ing no  oxygen,  and  potassium  perchlorate,  containing  46.2% 

24 


COMMON  LABOEATOEY  METHOD.        25 

oxygen,  are  formed  ;  and  some  oxygen  is  liberated.  For  the 
decomposition  of  the  potassium  perchlorate,  a  much  higher 
temperature  is  required.  The  final  result  is  that  the  potassium 
chlorate  has  been  broken  up  completely  into  potassium  chloride 
and  oxygen.  Potassium  chloride  is  not  volatile  at  the  temper- 
ature used,  and  therefore  remains  behind  in  the  retort. 

20.  Common  Laboratory  Method. — The  evolution 
of  oxygen  takes  place  much  more  easily,  and  at  about 
200°  C.,  if  we  heat  a  mixture  of  potassium  chlorate 
with  manganese  dioxide,  or  ferric  oxide,  instead  of 
potassium  chlorate  alone.  This  is  the  common  labora- 
tory method. 

Operation. 

Approximately  equal  parts  by  weight  of  manganese  dioxide 
and  potassium  chlorate  are  powdered  separately  in  clean  mor- 
tars, and  then  mixed  carefully  on  clean,  sized  paper.  Before 
decomposing  the  whole  mixture,  we  test  its  quality  by  heat- 
ing a  small  portion  in  an  open  test  tube.  If  there  is  evidence 
of  violent  combustion,  or  if  large  sparks  appear  in  the  test 
tube,  we  reject  the  mixture  and  make  a  fresh  one.  A  few 
small  sparks  indicate  only  traces  of  impurity  (dust,  etc.).  If 
the  mixture  is  found  sufficiently  pure  it  is  placed  in  a  large  test 
tube,  or  a  small  flask  (they  may  be  of  ordinary  soft  glass),  pro- 
vided with  a  delivery  tube  terminating  under  water.  Oxygen 
comes  off  readily  when  a  gentle  heat  is  applied  to  the  flask. 

If  the  materials  used  are  not  of  "chemically  pure"  (c.  p.) 
grade,  the  oxygen  will  contain  impurities  ;  these  may  be  re- 
moved sufficiently  by  allowing  the  gas  to  bubble  through  sodium 
hydroxide  solution. 

In  the  decomposition  of  potassium  chlorate,  as  in  the 
reactions  studied  in  Chapter  I,  the  relation  between  the 


26  OXYGEN. 

quantity  -of  material  taken  and  that  of  each  of  the  prod- 
ucts formed  is  definite  and  constant. 

Thus,  122.5  grams  of  potassium  chlorate  give  48  grams  of 
oxygen  and  74.5  grains  potassium  chloride. 

The  ease  with  which  potassium  chlorate  decomposes 
in  the  presence  of  manganese  dioxide  was  unexplained 
for  a  long*  time  ;  but  it  is  probably  due  to  the  fact  that 
the  two  substances  react  to  form  intermediate  com- 
pounds, which  are  decomposed  again.  The  heating  of 
a  mixture  of  manganese  dioxide  and  potassium  chlorate 
to  about  200°  to  250°  C.  thus  results  in  the  decompo- 
sition of  the  potassium  chlorate,  while  it  leaves  the  man- 
ganese dioxide  unchanged. 

21.  Other  Methods  of  Preparing  Oxygen.  —  Three 
other  methods  of  preparing  oxygen  will  be  described 
briefly;  these  are, — 

(1)  Decomposition  of  mercuric  oxide. 

(2)  Decomposition  of  barium  peroxide. 

(3)  Decomposition  of  manganese  dioxide. 

The  first  method  is  the  historic  one  of  Priestley,  who 
discovered  oxygen  in  1771,  and  of  Scheele  (pronounced 
Shala),  who  discovered  it  independently  in  1774. 

When  mercury  is  heated  in  air  to  a  temperature  a 
little  below  the  boiling  point  of  the  mercury  (mercury 
boils  at  357°  C.),  it  unites  .with  a  definite  weight  of 
oxygen  to  form  mercuric  oxide,  or  red  oxide  of  mercury. 


PHYSICAL  PROPERTIES   OF   OXYGEN.  27 

25  parts  by  weight  of  mercury  and  2  parts  of  oxygen  give  27 
parts  of  mercuric  oxide. 

Mercuric  oxide  is  a  stable  compound  at  ordinary  tem- 
eratures  ;  but  when  it  is  heated  to  a  temperature  a  little 
higher  than  that  at  which  it  was  formed,  it  is  decomposed 
again  into  mercury  and  oxygen. 

27  grams  of  mercuric  oxide  always  yield  25  grams  of  mercury 
and  2  grams  of  oxygen. 

Barium  peroxide,  the  second  substance  named  above  as  a 
source  of  oxygen,  is  a  white  solid  which  gives  up  half  of  its 
oxygen  when  heated  ;  the  other  half  remains  in  combination 
with  the  barium  in  the  compound  barium  monoxide.  Under 
appropriate  conditions  barium  monoxide  takes  up  from  the  air 
as  much  oxygen  as  it  already  holds,  and  thus  forms  the  per- 
oxide. Apparatus  has  been  devised  in  which  these  changes 
take  place  alternately,  and  large  quantities  of  oxygen  are  thus 
produced  for  sale. 

169  grams  of  barium  peroxide  give  153  grams  of  barium 
monoxide  and  16  grams  of  oxygen. 

It  is  evident  that  the  oxygen  formed  in  both  the  first  and 
the  second  methods  is  taken  from  the  air. 

Manganese  dioxide,  the  third  substance  named  above,  is 
called,  also,  "  black  oxide  of  manganese.'1''  It  is  found  in  nature 
as  the  mineral  pyrolusite.  Manganese  dioxide  decomposes  at 
about  600°  C.,  giving  off  a  third  of  its  oxygen. 

261  grams  of  manganese  dioxide  give  229  grams  of  manga- 
nous-manganic  oxide  and  32  grams  of  oxygen. 

22.  Physical  Properties  of  Oxygen.  —  In  whatever 
way  it  is  prepared,  oxygen  is  colorless,  odorless,  and 
tasteless,  if  pure.  It  is  somewhat  heavier  than  air,  and 


28  OXYGEN. 

sixteen  times  as  heavy  as  hydrogen.  One  liter  of  oxygen 
at  0°  C.  and  760  mm.  pressure,  weighs  1.43  grams. 

Gaseous  oxygen  may  be  condensed  at  — 118°  C.  and 
50  atmospheres  (=  50  X  760  mm.)  pressure,  to  a  bright 
blue  liquid. 

Oxygen  is  more  than  twice  as  soluble  in  water  as 
hydrogen ;  100  c.c.  of  water  dissolve  about  4  c.c.  oxygen 
under  ordinary  conditions. 

Oxygen  is  the  most  abundant  element,  composing  about  half 
of  the  earth's  solid  crust,  eight-ninths  of  the  water,  and  23% 
by  weight  of  the  atmosphere.  It  is  an  essential  constituent  of 
all  living  things. 

23.  Chemical  Properties.  —  The  chief  chemical  prop- 
erty of  oxygen  is  its  energetic  support  of  combustion  ; 
for  substances  that  burn  in  air  burn  much  more  rapidly 
in  oxygen.  Thus,  a  pine  splinter  which  is  merely  glow- 
ing in  the  air  will  burst  into  flame  if  put  into  oxygen. 
Combustion  in  air  is  more  slow  than  in  oxygen,  because 
the  oxygen  of  the  air  is  diluted  with  almost  four  times 
its  volume  (=  more  than  three  times  its  weight)  of 
inert  gases  which  do  not  support  ordinary  burning 
at  all. 

Examples  of  substances  which  burn  readily  in  oxygen 
are :  Iron,  which  burns  with  scintillation,  forming  the 
magnetic  oxide  of  iron;  magnesium,  phosphorus,  and 
sulphur,  which  burn  with  intensely  brilliant  flames ;  and 
charcoal,  which  burns  with  a  yloiv,  as  in  air,  but  much 


OXIDES.  29 

more  brightly.     These  substances  unite  with  oxygen  in 
the  following  proportions  :  - 

21  grams  of  iron  and  8  grams  of  oxygen  give  29  grams  of 
magnetic  iron  oxide. 

3  grains  of  magnesium  and  2  grams  of  oxygen  give  5  grams 
of  magnesium  oxide. 

31  grams  of  phosphorus  and  40  grams  of  oxygen  give  71 
grams  of  phosphorus  pentoxide. 

1  gram  of  sulphur  and  1  gram  of  oxygen  give  2  grams  of 
sulphur  dioxide. 

3  grams  of  carbon  (charcoal)  and  8  grams  of  oxygen  give  11 
grams  of  carbon  dioxide. 

The  magnetic  oxide  of  iron  is  a  black  solid  ;  magne- 

sium oxide  and  phosphorus  pentoxide  are  white  solids 

-  the  latter  is  very  soluble  in  water  ;  sulphur  dioxide 

and  carbon  dioxide  are  colorless  gases.     Sulphur  dioxide 

has  the  characteristic  odor  of  burning  sulphur. 

24.  Oxides.  —  Because  of  the  readiness  with  which 
oxygen  unites  with  other  elements,  the  most  common 
compounds  in  which  elementary  bodies  are  found  are 
oxides.  Fluorine  and  the  argon  family  alone,  of  all  the 
elements,  form  no  oxygen  compounds  so  far  as  known. 


}7  oxides  maybe  made  directly  from  the  elements,  e.  g., 
water,  and  copper  oxide  ;  but  many  must  be  made  indirectly, 
e.  g.,  platinum  oxide.  To  obtain  the  latter  substance  the  metal 
platinum  must  first  be  converted  into  other  compounds,  and 
these  into  the  oxide.  Nitrogen,  too,  although  incombustible  in 
the  ordinary  sense,  yet  forms,  by  indirect  methods,  five  dif- 
ferent oxides. 


30  OXYGEN. 

25.  Oxidation  and  Reduction.  —  To  the  union  of 
oxygen  with  other  substances  we  give  the  name  oxida- 
tion, and  to  a  substance  which  gives  up  some  of  its 
oxygen  to  another  body  the  name  oxidizing  agent.     Of 
course  the  oxidizing  agent  is  itself  reduced,  i.  e.,  loses 
oxygen,  when  it  oxidizes  another  substance  —  there  can 
be  no  oxidation  without  a  corresponding  reduction.     The 
substances  mentioned  in  this  chapter  as  sources  of  oxy- 
gen are  all  oxidizing  agents ;  carbon  and  hydrogen  (c/. 
§  12)  are  common  examples  of  reducing  agents. 

26.  Deflagration.  —  When  a  solid  or  liquid  combus- 
tible substance  is  mixed  with  a  solid  or  liquid  oxidizing 
agent,  and   the   temperature  is  raised   sufficiently,  the 
combustion  does  not  proceed  from  one  part  of  the  com^ 
bustible  to  another,  as  is  the  case  when  the  combustible 
burns  in  air;  on  the  contrary,  union  takes  place  almost 
instantaneously  through  the  whole  mixture,  just   as  it 
does  through  a  mixture  of  gaseous  hydrogen  and  oxy- 
gen.    This  rapid   union  of  combustible   and  oxidizing 
agent  is  called  deflagration. 

A  common  case  of  deflagration  is  that  of  gunpowder,  which 
is  a  mixture  of  charcoal  and  sulphur  (reducing  agents)  with 
potassium  nitrate,  or  potassium  chlorate  (oxidizing  agents). 
When  ignited  in  air,  gunpowder  deflagrates.  In  an  enclosed 
space  the  same  action  takes  place,  but  the  gaseous  products  of 
the  combustion  are  held  for  an  instant  under  great  pressure. 
When  this  pressure  is  released,  the  expanding  gases  are 
capable  of  hurling  a  projectile,  or  of  tearing  apart  masses  of 
rock. 


SLOW  COMBUSTION.  31 

27.  Combustion.  —  When  a  substance  unites  directly 
with  gaseous   oxygen  we  speak  of  the  oxidation  as  a 
case  of  burning,  or  combustion.     Of   this  we  generally 
distinguish  two  kinds,  (1)  ordinary  and  (2)  slow  com- 
bustion. 

In  ordinary  combustion,  heat  is  produced  by  the  union 
of  combustible  with  oxygen  much  more  rapidly  than  it 
can  be  dispersed  by  conduction,  radiation,  etc.  ;  conse- 
quently the  temperature  of  the  burning  body  rises  far 
above  that  of  the  surrounding  medium.  Usually  a  part 
of  the  combustible,  or  of  the  products  of  combustion, 
becomes  incandescent ;  and  some  of  the  energy  liberated 
appears  in  the  form  of  light. 

28.  Slow  Combustion.  —  A  slow  combustion  occurs 
when   oxidation  takes  place  through  a  long  period  of 
time,  and,  therefore,  without  a  decided  rise  of  temper- 
ature.    This  can  occur  only  when  the  heat  is  dispersed 
as  rapidly  as  it  is  evolved. 

Decay,  e.g.,  of  wood,  is  a  form  of  slow  combustion  ;  so  is 
the  rusting  of  iron. 

The  temperature  of  the  bodies  of  animals  is  kept  up  by  the 
slow  oxidations  taking  place  within  them. 

Although  there  is  a  great  difference  in  temperature 
between  a  body  oxidizing  slowly  and  one  burning  in  the 
ordinary  way,  yet  the  amount  of  energy  actually 
evolved  by  the  oxidation  of  a  given  weight  of  a  sub- 
stance is  the  same  in  the  one  case  as  in  the  other. 

-i 


32  OXYGEN. 

Thus,  a  piece  of  magnesium  oxidizes  slowly  in  moist  air,  at 
the  ordinary  temperature,  to  form  a  white  powder  containing 
magnesium  oxide  and  water  ;  but  we  have  every  reason  to 
believe  that  the  quantity  of  energy  set  free  during  this  slow 
formation  of  magnesium  oxide  is  just  as  great  as  that  evolved 
when  an  equal  mass  of  the  metal  magnesium  burns  brightly 
in  the  air.  That  there  is  no  perceptible  heat  and  light  in  the 
former  case  is  due  to  the  fact  that  the  evolution  of  heat  is 
equaled  by  its  dispersion. 

29.  Spontaneous    Combustion.  —  When   the    evolu- 
tion of  heat  is  only  a  little  in  excess  of  its   dispersion, 
the  combustion  is  apparently  a  slow  one ;  after  a  while, 
however,  enough  heat  accumulates  to  set  the  body  on 
fire.     Slow  oxidation  explains  so-called-  "  spontaneous  " 
combustions,  by  which  heaps  of  oily  rags,  etc.,  ignite 
without  apparent  cause. 

Spontaneous  combustion  may  be  illustrated  by  means  of  a 
solution  of  phosphorus  (only  a  small  amount  must  be  used) 
in  carbon  disulphide.  If  this  solution  is  poured  upon  a  tilter 
paper  supported  on  a  ring  stand,  the  phosphorus  will  soon  take 
fire  "spontaneously."  The  explanation  of  the  phenomenon 
is  that  the  evaporation  of  the  carbon  disulphide  leaves  the 
phosphorus  in  the  pores  of  the  paper,  where  it  oxidizes  ;  and 
the  heat  generated,  being  prevented  from  escaping  by  the  non- 
conducting filter  paper,  soon  raises  the  temperature  of  some 
part  of  the  paper  to  the  ignition  temperature  of  the  phos- 
phorus, 

30.  Ignition    Temperature.  —  In   all  ordinary  com- 
bustion   there    is    a    definite    temperature,    called    the 
ignition  temperature,    or   the    kindling   temperature,    to 


COMBUSTION  IN  AIR;   DRAFTS.  33 

which  the  combustible  substance  must  be  heated  in 
order  that  it  may  begin  to  unite  with  the  gas  supporting 
the  combustion.  The  burning  substance  must  not  only 
be  heated  up  to  the  kindling  temperature,  but  it  must 
be  kept  at  least  as  liigli  as  this  temperature,  or  combus- 
tion will  cease. 

The  ignition  temperature  is  different  for  different  sub- 
stances. Thus,  ordinary  phosphorus  bursts  into  flame,  in 
air,  at  about  40°  C.  ;  while  for  sulphur  the  kindling  tempera- 
ture is  about  260°  C.  In  some  cases  the  temperature  of 
ignition  is  far  below  the  ordinary  temperature. 

The  heat  evolved  by  the  burning  of  one  part  of  a 
substance  serves  to  raise  other  parts  to  the  ignition 
point;  illustrations  of  this  are  found  in  the  burning  of 
wood,  paper,  etc.  In  the  ordinary  match  the  ignition 
temperature  of  the  material  composing  the  head  is 
reached  by  friction;  and  the  heat  generated  by  the 
combustion  of  the  head  serves  to  raise  the  wood  of  the 
match  to  its  ignition  temperature. 

31.  Combustion  in  Air;  Drafts.  —  The  heat  given 
off  in  combustion  is  taken  up,  not  only  by  new  portions 
of  the  burning  body,  but  also  by  fresh  portions  of  the 
surrounding  gaseous  medium ;  if,  therefore,  the  latter 
is  diluted,  some  of  the  heat  available  will  be  used  in 
raising  the  temperature  of  the  diluting  gas,  as  well'  as 
that  of  the  gas  taking  part  in  the  combustion.  Hence, 
in  the  air,  which  is  a  mixture  of  oxygen  and  nitrogen 


34  OXYGEN. 

(chiefly),  combustion  is  much  slower  than  in  pure  oxy- 
gen, since  the  nitrogen,  although  it  takes  no  part  in  the 
combustion,  yet  takes  up  much  of  the  heat  evolved  by 
the  combustion. 

Moreover,  when  the  products  of  the  combustion  are 
gaseous,  they  dilute  the  oxygen  still  further.  Thus, 
when  charcoal  burns  in  an  enclosed  portion  of  air,  com- 
bustion ceases  long  before  all  the  oxygen  is  exhausted, 
for  the  reason  that  the  carbon  dioxide  gas  which  is 
formed  dilutes  the  oxygen. 

If  the  products  of  the  combustion  are  removed  as  rapidly 
as  formed,  the  combustion  will  be  much  more  nearly  com- 
plete. Thus,  when  phosphorus  burns  in  an  enclosed  portion 
of  air  over  water,  the  product  of  the  combustion  —  phosphorus 
pentoxide  —  dissolves  in  the  water,  and  does  not  dilute  the 
oxygen ;  as  a  result  the  oxygen  is  practically  all  taken  up 
by  the  phosphorus.  For  the  same  reason  the  combustion  of 
charcoal  in  an  enclosed  portion  of  air  may  be  made  much 
more  nearly  complete  if  carried  out  over  sodium  hydroxide 
solution. 

The  removal  of  the  products  of  combustion  from  the 
"  sphere  of  action  "  is  accomplished  in  ordinary  burning 
by  means  of  drafts,  which  also  bring  fresh  supplies  of 
air.  A  moderate  draft  is  thus  beneficial  to  combustion. 
The  air  current  may,  however,  have  such  a  velocity  that 
the  heat  evolved  in  the  combustion  is  not  sufficient  to 
raise  the  temperature  of  the  air  supplied  and  of  fresh 
portions  of  the  burning  body  to  the  kindling  tempera- 


FLAMES. 


35 


cure.     Hence    combustion  ceases.     A  flame  may    thus 
be  "  blown  out "  by  a  strong  current  of  air. 

32.  The  Safety  Lamp. — The  lowering  of  the  tem- 
perature of  a  flame  below  the  kindling  temperature  is 
admirably   illustrated    in    the    safety 

lamp  devised  by  Sir  Humphry  Davy. 
The  lamp  consists  of  an  ordinary 
lantern  entirely  surrounded  by  wire 
gauze.  When  such  a  lamp  (Fig.  12) 
is  carried  into  an  explosive  mixture 
of  gases,  e.g.,  hydrogen  and  air,  the 
gases  diffuse  through  the  wire  gauze 
and  burn  inside  the  lamp ;  but  the 
heat  generated  is  conducted  away 
by  the  wire  gauze  instead  of  being 
communicated  to  the  explosive  mix- 
ture outside ;  hence  an  explosion 
of  the  gases  outside  of  the  lamp  is 
avoided. 

Safety  lamps  are  used  to  prevent  the 
explosion  of  the  "  fire-damp,"  —  a  mix- 
ture of  marsh  gas  and  air,  —  which  often 
occurs  in  mines.  FIG.  12. 

33.  Flames.  —  Aflame  is  a  gas  in  combustion.     To 
burn  with  a  flame,  a  substance  must  either  be  gaseous 
itself,  or  it  must  evolve  gaseous  products.     Such  sub- 
stances as  magnesium,  sulphur,  phosphorus,   and  wax 


36  OXYGEN. 

burn  with  flame  because  they  are  first  converted  into 
the  gaseous  form;  wood  and  soft  coal,  because  they 
evolve  combustible  gaseous  products ;  but  charcoal, 
which  contains  practically  110  volatile  constituents, 
merely  glows.  The  structure  of  flames  will  be  taken 
up  later. 

34.  Reversed  Combustion.  —  We  have  spoken  of 
combustion  heretofore  as  the  union  of  the  burning  body 
with  oxygen ;  other  gases,  however,  may  be  supporters  of 
combustion  just  as  truly  as  oxygen. 

Thus,  a  jet  of  burning  hydrogen  continues  to  burn  in  bro- 
mine vapor ;  and  phosphorus  burns  in  chlorine  much  as  in 
oxygen. 

If  both  the  combustible  and  the  supporter  of  its  coin, 
bustion  are  gaseous,  the  combustion  may  be  reversed. 
Thus,  oxygen  may  become  the  burning  body  and  illumi- 
nating gas    the    supporter  of   combustion. 
This   reversal  may  be    shown    by   a   very 
simple  experiment. 

A  bottle  (Fig.  13)  is  supported,  mouth  do\vn- 
ward,  and  filled  with  illuminating  gas  by  dis- 
jL-  /xC  placing  the  air.     The  gas  at  the  mouth  of  the 

bottle  is  then  lighted,  and  while  it   is  burning 
^        """       a  jet  of  oxygen  is  brought  up  into  the  bottle. 
FIG.  13.         The   oxygen  takes  fire   at   the   bottle's   mouth 
and    burns    in    the    atmosphere    of    illuminat- 
ing gas.     The  oxygen  jet  of  the  preceding  experiment  may 


EXERCISES.  37 

be  replaced  by  a  deflagrating  spoon  of  potassium  chlorate  which 
has  been  heated  so  that  it  gives  off  oxygen. 

35.  Exercises. 

1.  How  many  grams  of  mercury  will  b^  formed  by  the  dej 
composition   of   43.2   grams  of  mercuric  oxide  ?     How  many 
grams  of  oxygen  ? 

2.  How  many  grams  of  the  magnetic  oxide  of  iron  will  be 
formed  when  50  grams  of  iron  burn  in  oxygen  ? 

3.  How  many  grams  of  manganese  dioxide  are  needed  to 
give,  when  decomposed  by  heat,  12  grams  of  oxygen?     How 
much  manganous-manganic  oxide  is  formed  at  the  same  time  ? 

4.  How  much  magnesium  is  contained  in  30  grams  of  mag- 
nesium oxide  ?     What  per  cent  of  magnesium  oxide  is  mag- 
nesium ?     Oxygen  ? 

5.  Calculate  the  per  cent  of  oxygen   in   phosphorus   pen- 
toxide. 

6.  How  many  c.c.  of  oxygen  can  be  made  from  1.2  grams  of 
potassium   chlorate    when   1    c.c.    of   the   gas   weighs    0.0014 
grams  ? 

7.  How  many  grams  of  potassium  chlorate  must  be  decom- 
posed to  fill  a  36-liter  gasometer  (a  vessel  for  storing  gases) 
with-  oxygen  at  a  temperature  and  a  pressure  at  which  1  liter  of 
oxygen  weighs  1.25  grams  ? 

8.  What  is  the  weight  of  the  carbon  dioxide  formed  by  the 
combustion  of  10  grams  of  carbon  in  oxygen  ?     What  will  be 
the  volume  of  the  carbon  dioxide  under  conditions  at  which 
one  liter  of  the  gas  weighs  2  grams  ? 


CHAPTER  III. 


WATER. 

36.  Nature  of  Water.  —  The  union  of  the  elements 
composing  water  is  so  strong  that  water  itself  was  be- 
lie vecl  to  be  an  element  until  1781.  In  that  year,  Cav- 
endish, who  had  discovered  hydrogen  in  1766,  succeeded 
in  synthesizing  (=  putting  together)  water  from  hydro- 
gen and  oxygen,  and  thus  proved  its  compound  nature. 

37.  Electrolysis  o  f 
Water.  —  The  fact  that 
water  is  a  compound  is 
proved  analytically  by  its 
decomposition  b  y  the 
electric  current.  T  h  e 
operation  is  carried  <out 
as  follows :  — 


FIG.  M. 


A  current  from  an  elec- 
tric battery  (Fig.  14)  —  that 
from  four  Gre'net  cells  in 

series  is  adequate  —  is  passed  between  platinum  electrodes 
through  water  containing  about  five  per  cent  of  its  weight  of 
sulphuric  acid.  While  the  current  is  passing,  bubbles  of  gas 
gather  upon  the  electrodes,  and  rise  from  them  through  the 
liquid.  The  gases  may  be  collected  separately  by  inverting 
over  each  electrode  a  tube  filled  with  the  dilute  acid.  The 


SYNTHESIS  OF   WATER,   BY   VOLUME.  39 

rates  at  which  the  two  gases  from  the  electrodes  collect  in  the 
tubes  is  not  the  same  ;  one  of  them — that  at  the  negative  ( — ) 
electrode  —  collects  a  little  more  than  twice  as  rapidly  as  the 
other.  This  electro-positive  (-)-)  gas  is  hydrogen  ;  that  at  the 
positive  electrode  is  oxygen. 

The  relation  between  the  volume  of  the  hydrogen  and 
that  of  the  oxygen  is  much  more  nearly  2  :  1  than  is 
usually  shown  by  this  experiment.  That  the  ratio  is 
generally  too  large  is  due  to  several  slight  errors,  one 
of  which  is  that  the  oxygen  is  much  more  soluble  in 
the  dilute  acid  than  the  hydrogen  This  error  may  be 
avoided  if  the  gases  are  not  collected  until  the  liquid 
has  become  saturated  with  them. 


The  electrolysis  of  dilute  sulphuric  acid  is  not  a  direct  der 
composition  of  water  by  the  electric  current;  for  pure  water  is 
scarcely,  if  at  all,  electrolyzed  in  the  absence  of  sulphuric  acid, 
or  some  similar  substance.  It  is  the  water,  however,  that  is 
actually  used  up;  hence  we  speak  of  the  result  as  an  elec- 
trolysis of  water.  The  function  of  the  sulphuric  acid  will  be 
discussed  later. 


38.  Synthesis  of  Water,  by  Volume. —  The  elec- 
trolysis of  water  gives  two  volumes  of  hydrogen  for 
every  volume  of  oxygen;  but  the  complete  proof  that 
this  is  the  proportion  in  which  these  gases  are  united 
in  water  follows  from  the  volumetric  synthesis  of  water. 
Thus  an  additional  fact  may  be  learned,  viz.,  that  the 
relation  between  the  volume  of  the  water  (steam)  pro- 


40 


WATER. 


duced   and   the    volume   of  the   oxygen   and  hydrogen 
taken  is  a  simple  one. 

The    apparatus    required    to  demonstrate    the    volu- 
metric synthesis  of  water  is  shown  in  Fig.  15. 


Coil 


A  graduated  tube  closed 
at  one  end  has  two  plati- 
num wires  passing  through 
the  walls  and  almost  meet- 
ing at  the  closed  end.  This 
tube  —  it  is  called  a  "  eudi- 
ometer "  tube  —  is  filled 
with  mercury,  arid  its  open 
end  is  placed  under  mer- 
cury. Part  of  the  mercury 
in  the  eudiometer  is  then 
displaced  by  a  mixture  of 
pure  hydrogen  and  oxygen 
put  together  in  the  propor- 
tions in  which  they  are 
obtained  by  electrolysis . 
The  apparatus  shown  in 
Fig.  16  is  very  convenient 
for  preparing  the  oxyhydro- 
gen  mixture.  The  eudi- 
ometer is  now  attached 
securely  to  a  rubber  tube 
containing  mercury,  and  is 

thus  put  into  communication  with  .a  leveling  tube  partly  full 
of  the  same  metal.  By  raising  or  lowering  the  leveler  the 
experimenter  can  compress  or  expand  the  gas  in  the  eudi- 
ometer. He  can  thus  get  the  volume  of  the  gas  at  atmospheric 
pressure  by  bringing  the  surface  of  the  mercury  in  the  eudi- 


Steam 


FIG.  15. 


SYNTHESIS   OF   WAT  EH,   BY    VOLUME. 


41 


ometer  and  that  in  the  leveler  into  the  same  horizontal 
plane.  When  the  connections  between  the  eudiometer  and 
the  leveler  have  been  made,  a  jacket  is  placed  about 
the  eudiometer,  the  platinum  wires  are  connected  with  a 
Ruhmkorff  coil,  and  steam 
is  passed  through  the  jacket 
until  the  volume  of  the  gas 
in  the  eudiometer  becomes 
constant.  This  volume  is 
read  accurately  at  atmos- 
pheric pressure.  The  leveler 
is  then  lowered,  so  as  to  put 
the  gas  under  diminished 
pressure;  the  spark  is  passed 
through  the  mixture;  and 
union  is  effected.  With  the 
steam  still  running  through 
the  jacket,  the  mercury  level 
in  the  eudiometer  and  in  the 

leveler  is  made  the  same;  and  the  volume  of  gas  in  the  tube 
is  read. 


The  gaseous  substance  in  the  tube  after  the  explosion 
is  steam ;  and  its  volume  is  only  two  thirds  as  great  as 
that  of  the  original  gases  ;  we  have  thus  proved  that  two 
volumes  of  hydrogen  and  one  of  oxygen  unite  to  produce 
two  of  steam. 

This  experiment  has  been  performed  many  times  and 
with  great  care, —  the  first  time  by  Humboldt  and  Gay- 
Lussac  in  1805, —  and  proves  conclusively  that  the  pro- 
portion by  volume  in  which  hydrogen  and  oxygen 


42  WATER. 

combine   to  form  water  is  the   same  as  that  in  which 
these  gases  are  obtained  from  water. 

If  the  apparatus  for  effecting  the  volumetric  synthesis  of 
water  is  used  without  the  steam  jacket,  the  proportions  of  the 
combining  gases  may  be  determined,  but  not  the  volume  of 
steam  produced.  The  volume  of  liquid  water  produced  by  the 
condensation  of  the  steam  is,  of  course,  very  small. 

39.  Synthesis  of  Water,  by  Weight.  —  A  knowl- 
edge of  the  exact  proportions,  by  weight,  in  which 
hydrogen  and  oxygen  are  combined  in  water  is  of 
such  importance  to  Chemistry  that  many  methods  have 
been  devised  for  the  determination  of  the  ratio.  The 
methods  are,  in  general,  of  two  classes. 

In  methods  of  the  first  class  a  known  weight  of 
hydrogen  is  passed  over  some  oxidizing  agent,  e.  g., 
cupric  oxide ;  the  hydrogen  is  thus  converted  into 
water,  which  is  collected  and  weighed.  The  gain  in 
weight  is,  evidently,  oxygen. 

In  methods  of  the  second  class  a  known  weight  of 
the  oxidizing  agent  is  reduced  in  a  stream  of  hydrogen, 
and  the  weight  of  the  water  formed  is  determined. 
Here  the  loss  in  weight  of  the  oxidizing  agent  is  plainly 
equal  to  the  weight  of  the  oxygen  which  united  with 
the  hydrogen  ;  for  the  water  formed  contains  the  oxygen 
lost  by  the  oxidizing  agent. 

The  apparatus  for  one  method  of  the  second  class  is 
shown  in  Fig.  17. 


NATURAL    WATER  AND  ITS  IMPURITIES. 


43 


Hydrogen,  purified  and  dried  by  alkaline  permanganate  solu- 
tion (B)  and  calcium  chloride  (C)  respectively,  is  passed  over 
heated  copper  oxide  contained  in  a  porcelain  boat  (D) ,  and  the 
water  produced  is  collected  in  all-tube  of  calcium  chloride  (E). 
A  guard  tube  of  calcium  chloride  (F)  excludes  the  water  of 
the  air. 

The  cupric  oxide  is  weighed  before  and  after  the  ex- 
periment; its  loss  in  weight  is  oxygen.  The  calcium 
chloride  tube,  too,  is  weighed  before  and  after  the 
experiment ;  its  gain  in  weight  represents  water. 

Berzelius  and  Dulong  carried  out  this  experiment,  in 
1819,  with  the  following  results :  — • 


FIG.  17. 

"Weight  of  water  taken  up  by  calcium  chloride    =  30.519  g. 
Loss  in  weight  of  the  cupric  oxide  (=  oxygen)    =  27.1^9  g. 

Therefore,  weight  of  hydrogen  united  with  > 
27.129  g.  oxygen  j 

27.129  :  3.39  ::  x  :  1  ;   whence  x  =  8.002  -f  ,    the   ratio  of 
oxygen  to  hydrogen  in  water. 

40.  Natural  Water  and  Its  Impurities.  — The  water 
that  falls  upon  the  earth's  land  surface  gets  back  to  the 


44  WATER. 

sea  in  various  ways,  but  rarely  without  leaching  out 
soluble  substances  from  the  soil.  Natural  water  there- 
fore contains  more  or  less  impurity.  The  character  of 
the  impurity  depends,  (1)  upon  the  substances  present 
in  the  air  through  which  the  water  fell  to  the  earth ; 
(2)  upon  the  soil  through  or  over  which  the  water  has 
flowed ;  and,  also,  (3)  upon  the  opportunities  the  water 
has  had  of  losing  material  previously  gathered.  Even 
rain  water  is  far  from  pure,  for  it  gathers  much  dust, 
both  organic  and  inorganic,  and  many  gaseous  impurities 
of  the  air,  e.  g.,  ammonia. 

Water  which  penetrates  the  earth's  surface  usually 
finds  soluble  substances,  both  solid  and  gaseous.  The 
most  common  soluble  solids  found  in  water  are,  prob- 
ably, common  salt,  magnesium  chloride,  and  gypsum;  of 
the  gases,  carbon  dioxide  and  hydrogen  sulphide.  Water 
charged  with  carbon  dioxide  has  the  power  to  dissolve 
limestone  ;  hence  this  substance  is  a  common  ingredient 
of  natural  water,  even  of  moderately  "soft"  water,  as 
is  proved  by  the  incrustations  of  limestone  in  vessels  in 
which  such  water  is  habitually  heated. 

Water  charged  with  hydrogen  sulphide  is  called  sul- 
phur water. 

The  water  which  flows  over  the  earth  as  rivers  gathers 
its  peculiar  organic  impurities  from  the  land.  These  may 
consist  of  micro-organisms  washed  down  by  surface  water, 
or,  in  the  case  of  rivers  passing  large  cities,  of  sewage.  If 
the  river  is  sluggish,  these  impurities  are  not  easily  removed ; 


THE  PURIFICATION  OF   WATER.  45 

but  if  it  has  a  rapid  current,  and  especially  if  there  are 
rapids  and  waterfalls  in  its  course,  the  river  soon  purifies 
itself  by  bringing  its  impurities  into  contact  with  the  oxygen 
of  the  air,  which  destroys  them. 

41 .  Sea  Water.  —  Since  the  sea  is  the  ultimate  desti- 
nation of  most  of  the  water  that  falls  upon  the  land,  it 
is  evident   that  the  material  dissolved  by  fresh  water 
will  accumulate  in  the  ocean.     Indeed,  about  four  per 
cent  of  sea  water  consists  of  dissolved   material,  three- 
fourths  of  which  is  common  salt.     It  is  probable  that 
greater  or  smaller  amounts  of  all  the  substances  com- 
posing the  crust  of  the  earth  may  be  found  in  the  sea. 

42.  The  Purification  of  Water.  —  Water  may  usually 
be  purified  by  filtration  or  by  distillation. 

Filtration  serves  not  only  to  remove  insoluble  sub- 
stances, but  also  to  oxidize  many  organic  impurities  by 
bringing  them  into  intimate  contact  with  air. 

When  water  is  raised  to  the  boiling  temperature,  most  of  the 
micro-organisms  contained  in  the  water  are  killed,  and  at  least 
one  of  its  inorganic  impurities,  viz.,  calcium  carbonate  (lime- 
stone), is  rendered  insoluble.  Soluble  impurities,  however, 
still  remain.  To  get  Avater  free  from  these  it  must  be  distilled. 

Distillation  consists  in  converting  a  liquid  into  vapor, 
and  then  condensing  the  vapor  to  the  liquid  state. 
When  water  is  distilled,  all  impurities  more  volatile 
than  the  water  will  appear  in  the  first  portions  of  the 


46 


WATER. 


distillate ;  all  much  less  volatile  will  remain  behind  in 
the  retort. 

The  usual  form  of  distilling  apparatus  used  in  labora- 
tories is  shown  in  Fig.  18. 

The  condenser  in  the  figure  is  called  a  Liebig's  condenser. 
It  consists  of  an  inner  tube  through  which  the  evolved  vapor 
is  passed  for  condensation,  and  of  an  outer  jacket  through 
which  a  stream  of  cold  water  is  kept  running  in  the  direction 
shown  by  the  arrows. 


FIG.  18. 


Although  water  can  be  obtained  reasonably  pure  by 
distillation,  chemically  pure  water  is  very  difficult  to 
prepare.  Even  if  water  is  pure  when  freshly  distilled, 
as  shown  by  its  leaving  no  residue  when  evaporated  in 
a  platinum  dish,  it  cannot  be  kept  pure  long,  owing  to 


PROPERTIES  OF   WATER.  47 

its  tendency  to  act  upon  the  glass  or  porcelain  vessels 
in  which  it  is  stored  or  used. 

Distilled  water  is  "  flat "  to  the  taste.  This  is  due  largely 
to  the  fact  that  distilled  water  has  lost  the  gases  present  in 
natural  water.  Distilled  water  may  therefore  be  made  much 
more  palatable  by  shaking  it  thoroughly  with  air. 

43.  Hard  and  Soft  Water.  —  Water  which  contains 
much  gypsum,  limestone,  or  similar  substances  in  solu- 
tion does  not  wet  the  skin  readily,  and  is  therefore  called 
"hard"  water.     When  soap  is  put  into  such  water  it 
does  not  dissolve  readily,  but  forms  an  insoluble  scum. 
It  is  only  after  the  separation  of  this  scum  that  soap 
will  dissolve  in  quantity  and  form  permanent  suds.     As 
will  be  explained  later,  the  hardness  of  water  containing 
only  limestone  is  temporary  because  it  may  be  removed 
by  boiling ;  if  gypsum  is  present,  however,  the  water  is 
permanently  hard  and  can  be  "  softened  "  only  by  the  use 
of  washing  powders,  etc.,  which  are  capable  of  convert- 
ing the  gypsum  into  insoluble  forms. 

44.  Properties  of  Water.  —  Pure  water  is  practically 
odorless  and  tasteless.     In  small  quantities  it  has  no 
color,  but  in  large  masses  it  is  blue. 

The  specific  heat  of  water  is  high,  more  heat  being 
required  to  raise  the  temperature  of  a  given  weight  of 
water  one  degree  than  is  required  in  the  case  of  any 
other  substance  except  hydrogen. 


48  WATEtt. 

The  latent  heat  both  of  water  and  of  steam  is  very 
great.  When  a  given  weight  of  water  at  0°  C.  is  frozen 
to  ice  at  0°  C.,  it  gives  off  enough  heat  to  raise  the 
temperature  of  an  equal  weight  of  water  from  0°  C.  to 
80°  C.  When  a  given  weight  of  steam  at  100°  C.  con- 
denses to  water  at  100°  C.,  the  heat  evolved  is  sufficient 
to  raise  the  temperature  of  about  5.37  times  the  weight 
of  water  from  0°  C.  to  100°  C. 

The  boiling  point  of  water  is  100°  C.  at  760  mm. 
pressure.  Since  the  boiling  point  of  a  substance  is  the 
temperature  at  which  the  pressure  of  its  vapor  just  ex- 
ceeds that  of  the  atmosphere,  the  "  vapor  tension  "  of 
water  at  100°  C.  must  be  760  mm.  At  21  mm.  pres- 
sure, water  boils  at  23°  C. ;  at  3,581  mm.,  at  150°  C. 

The  freezing  point  of  water  (=  melting  point  of  ice) 
is  0°  C.  at  760  mm.  pressure.  Water  expands  on  freez- 
ing, 10  c.c.  of  the  liquid  becoming  about  10.1  c.c.  of 
ice.  At  about  4°  C.,  water  is  at  its  maximum  den- 
sity. Its  relative  density  at  this  temperature  is  taken 
as  1. 

Water  is  a  poor  conductor  of  heat  and  of  the  electric 
current. 

45.  Steam  and  its  Dissociation.  —  Steam,  which  is 
water  in  the  condition  of  a  vapor,  is  nine  times  as  heavy 
as  hydrogen.  It  is  so  stable  that  it  does  not  begin  to 
decompose  into  its  elements  until  it  is  heated  to  about 
1000°  C.  Above  1000°  C.  the  amount  of  decomposition 
increases  with  the  temperature,  until  at  2500°  C.  about 


ACTION   OF  SODIUM   UPON   WATER.  49 

half  of  the  steam  is  no  longer  steam,  but  oxygen  and 
hydrogen  uncombined.  No  matter  how  long  steam  is 
kept  at  2500°  C.  it  cannot  be  decomposed  completely,  for 
the  reason  that  side  by  side  with  the  decomposition  of 
steam  into  hydrogen  and  oxygen  there  is  a  recombina- 
tion of  these  elements  to  form  steam.  At  every  tem- 
perature, therefore,  between  1000°  C.  and  the  (high) 
temperature  at  which  the  decomposition  of  steam  is 
complete,  a  condition  of  equilibrium  is  soon  reached,  at 
which  as  much  steam  is  produced  in  a  given  time  as  is 
decomposed  in  the  same  time.  Hence  the  change  pro- 
ceeds no  farther  unless  the  temperature  is  raised.  If 
the  temperature  is  lowered,  enough  hydrogen  and  oxy- 
gen recombine  to  produce  equilibrium  at  the  lower 
temperature. 

A  decomposition  like  that  of  steam  is  called  a  disso- 
ciation. 

46.  Action  of  Sodium  upon  Water.  —  In  the  first 
chapter  reference  was  made  to  the  action  of  sodium 
upon  water  as  a  means  of  preparing  hydrogen.  This 
reaction  will  now  be  considered  more  fully. 

Sodium  is  a  soft  solid,  somewhat  lighter  than  water. 
It  has  a  silvery  luster  when  freshly  cut,  but  tarnishes 
quickly  in  ordinary  air.  It  is,  therefore,  kept  under  lig- 
roin  or  kerosene.  Although  we  commonly  think  of  a 
metal  as  hard  and  heavy,  sodium,  which  has  neither  of 
these  properties,  is  yet  one  of  the  best  representatives 
of  the  class  of  metals.  This  is  due  to  its  chemical  prop- 


50  WATER. 

erties,  the  most  important  of  which,  for  our  present  pur- 
pose, is  its  behavior  toward  water. 

When  a  piece  of  sodium  is  thrown  upon  water  (this  is  done 
at  arm's  length  to  avoid  danger  from  spattering),  it  at  once 
attacks  the  water,  melts,  assumes  a  globular  form,  and  then 
swims  about  until  dissolved.  If  a  lighted  match  is  held  near 
the  sodium  while  it  is  floating  upon  water,  a  flame  will  appear ; 
this  is  burning  hydrogen.  The  hydrogen  may  be  collected  by 
placing  the  sodium  in  a  short  piece  (1  cm.  long)  of  glass  tubing, 
and  then  plunging  it  quickly  by  the  aid  of  tongs  under  the 
mouth  of  a  bottle,  or  a  test  tube,  filled  with  water  and  inverted 
in  a  pan  of  water. 

Water  in  which  a  sufficient  quantity  of  sodium  has 
dissolved  possesses  new  properties.  It  feels  soapy  to  the 
touch,  has  a  bitter  taste,  and  turns  many  vegetable  col- 
ors, e.  g.,  red  litmus  to  blue.  If  the  water  is  evaporated, 
a  white  substance  will  remain,  which  is  'sodium  hydrox- 
ide, or  caustic  soda.  It  is  this  substance  that  gives  the 
water  its  new  properties. 

47.  Quantitative  Study  of  the  Reaction ;  Hydrox- 
ides. —  Sodium  hydroxide  is,  as  its  name  indicates,  a 
compound  of  sodium,  hydrogen,  and  oxygen.  Since 
sodium  hydroxide  contains  hydrogen,  it  is  evident  that 
not  all  of  the  hydrogen  of  water  is  set  free  by  sodium. 
As  a  matter  of  fact,  only  one-half  of  the  hydrogen  is  so 
liberated,  the  remaining  half  being  in  combination  with 
sodium  and  oxygen.  This  appears  from  a  quantitative 
study  of  the  reaction. 


STUDY  OF  THE  REACTION;  HYDROXIDES.  51 

If  23  grams  of  sodium  had  been  placed  in  contact 
with  a  quantity  of  water  greater  than  18  grams,  it  would 
have  acted  upon  18  grams  of  water  only,  and  would 
have  formed  40  grams  of  sodium  hydroxide,  and,  at 
standard  temperature  and  pressure,  a  little  over  11  liters 
of  hydrogen.  The  hydrogen  formed  would  weigh  prac- 
tically 1  gram.  To  summarize  the  results  of  the  reac- 
tion, quantitatively  as  well  as  qualitatively :  — 


23  grams  of  sodium  and  18  grams  of  water  react  to  give  40 
grams  of  sodium  hydroxide  and  1  gram  of  hydrogen. 


As  we  have  already  learned,  18  grams  of  water  con- 
sist of  2  grams  of  hydrogen  and  16  of  oxygen.  Further- 
more, 40  grams  of  sodium  hydroxide  would  give,  when 
decomposed,  23  grams  of  sodium,  16  grams  of  oxygen, 
and  1  gram  of  hydrogen.  It  is  evident,  therefore,  that 
sodium  replaces  only  half  of  the  hydrogen  of  water  in 
forming  sodium  hydroxide. 

If  we  had  used  the  metal  potassium,  the  results  would 
have  been  similar,  viz.  :  — 


39  grams  of  potassium  and  18  grams  of  water  react  to  pro- 
duce 56  grams  of  potassium  hydroxide  and  1  gram  of  hydrogen. 


The  56  grams  of  potassium  hydroxide  consist  of  39 
grams  of  potassium,  16  grams  of  oxygen,  and  1  gram 
of  hydrogen.  Here  again  one-half  of  the  hydrogen  of 


52  WATER. 

the  water  decomposed  is  liberated ;  and  the  remaining 
half  is  retained  in  the  hydroxide.  In  fact,  the  hydrox- 
ides of  all  of  the  metals  may  be  considered  to  be  water 
with  half  of  its  hydrogen  replaced  by  a  metal. 

48.  The  Action  of  Metals  upon  Hydroxides.  —  By 

the  use  of  proper  methods,  the  hydrogen  of  sodium 
hydroxide  may  be  replaced  by  sodium,  and  that  of 
potassium  hydroxide  by  potassium.  The  resulting  sub- 
stances will  be  sodium  and  potassium  oxides.  The  re- 
actions take  place  as  follows  :  — 

h 

40  grams  of  sodium  hydroxide  and  23  grams  of  sodium  give 
62  grams  of  sodium  oxide  and  1  gram  of  hydrogen. 

Also,  56  grams  of  potassium  hydroxide  and  39  grams  of 
potassium  give  94  grams  of  potassium  oxide  and  1  gram  of 
hydrogen. 

To  replace  the  hydrogen  of  sodium  hydroxide  by 
sodium,  and  that  of  potassium  hydroxide  by  potassium 
is  a  somewhat  difficult  operation;  but  it  is  very  easy 
to  replace  the  hydrogen  of  these  hydroxides  by  alumi- 
num. The  resulting  substances  are  sodium-aluminum 
oxide  and  potassium-aluminum  oxide  *  respectively.  These 
compounds  are  similar  to  sodium  oxide  and  to  potas- 
sium oxide  in  that  they  are  water  with  all  of  its  hydro- 
gen replaced  —  in  two  stages  —  by  metallic  elements. 

All  of  the  facts  just  stated  are  given  in  the  following 
recapitulation :  — 

*  Cf .  aluminates,  §  424. 


WATER  MECHANICALLY  ENCLOSED. 


53 


HYDROGEN. 

OXYGEN. 

SODIUM. 

POTASSIUM. 

18  parts  by  weight  of  water  consist  of 

2 

16 

23 



40  parts  by  weight  of  sodium  hydroxide  consist  of 

1 

16 

62  parts  by  weight  of  sodium  oxide  consist  of 

1 

16 
16 

16 

46 

56  parts  by  weight  of  potassium  hydroxide  consist  of 

39 

94  parts  by  weight  of  potassium  oxide  consist  of 

78 

49.  Water  in  Combination.  —  Water  is  widely  dis- 
tributed, not   only  in   the  free   condition,  but  also    in 
combined   form.     Most  natural    substances   contain  it. 
This  is  true  not  only  of  animal  and  plant  tissues  and 
products,  but  even  of  inorganic  substances.     For  con- 
venience,1 we  may  distinguish  at  least    three    ways    in 
which  water  may  be  contained  in  other  substances :  — 

(1)  Mechanically  enclosed. 

(2)  As  "  water  of  crystallization." 

(3)  As  an  integral  part  of  the  substance. 

The  form  in  which  water  is  contained  in  a  substance  gener- 
ally appears  from  the  behavior  of  the  substance  when  heated. 

50.  Water  Mechanically  Enclosed.  —  Water  may  be 
held  mechanically  either  (1)  between  the  crystals  of  a 
substance,  or  (2)  in  its  pores.     In  either  case  the  water 


54  WATEE. 

is  given  off  when  the  substance  is  heated  gently. 
When  the  substance  contains  water  enclosed  between 
crystals,  however,  the  water  escapes  explosively ;  for 
the  crystal-mass  is  broken  in  pieces  by  the  steam  pro- 
duced. 

Such  substances  are  said  to  decrepitate. 

Illustrations  are  :  Common  salt  and  potassium  sulphate. 

51.  Water  of  Crystallization. — By  water  of  crys- 
tallization^ or  crystal-water,  we  mean  the  water  with 
which  some  substances  combine  when  they  crystallize 
from  aqueous  solution. 

When  substances  containing  water  of  crystallization 
are  heated,  they  usually  melt  while  the  water  escapes, 
and  then  assume  the  solid  form  again. 

The  loss  of  crystal-water  by  a  substance  is  accompanied  by 
a  loss  of  crystalline  structure  and  by  other  changes  in  proper- 
ties. Thus,  cupric  sulphate  is  a  white  solid,  but  blue  vitriol,  its 
ordinary  form,  is  cupric  sulphate  plus  water  of  crystallization. 
Crystallized  sodium  sulphate,  sodium  carbonate,  alum,  etc., 
all  contain  much  crystal-water. 

The  amount  of  crystal-water  which  Avill  combine  with 
a  given  weight  of  the  anhydrous  (i.e.,  water-free)  sub- 
stance is  definite  for  each  substance.  Thus,  90  grams 
of  water  are  united  with  159  grams  of  cupric  sulphate 
in  every  249  grams  of  blue  vitriol. 

The  color  of  a  substance  is  usually  the  same  when 


EFFLORESCENCE,   DELIQUESCENCE,   ETC.  55 

the  substance  is  combined  with  water  of  crystallization 
as  when  it  is  in  solution  in  water. 

By  no  means  all  crystalline  substances  contain  crystal- 
water.  Cane  sugar,  salt,  potassium  sulphate,  etc.,  crystallize 
from  their  solutions  in  water  without  taking  up  any  of  the 
water  as  crystal-water. 

52.  Water  an  Integral  Part  of  the  Substance.  — 

Many  substances  which  cannot  be  said  to  contain  water, 
yet  contain  hydrogen  and  oxygen  in  the  proportion 
in  which  these  elements  are  united  in  water.  Such 
compounds  are  the  substances  referred  to  as  having 
water  in  the  third  form  of  combination,  viz.,  as  an 
integral  part  of  the  substance.  Examples  are  sugar, 
starch,  cotton,  wood,  etc.  When  these  substances  are 
heated,  they  are  easily  decomposed,  liberating  water  and 
other  products,  while  a  residue  of  charcoal  remains 
behind. 

53.  Efflorescence,     Deliquescence,     Etc.  —  Certain 
substances   give  up  all  or   part    of  their   crystal-water 
when  exposed  to  the  air,  and  thus  lose  their  crystalline 
form.     Such  substances  are  said  to  effloresce. 

An  example  is  crystallized  sodium  carbonate,  which  be- 
comes a  non-crystalline  powder  when  exposed  to  the  air. 

Certain  other  substances,  on  the  contrary,  when  de- 
prived of  their  water  of  crystallization,  take  it  up 


56  WATER. 

again  by  absorbing  water  from  the  air  and  from  other 
substances.  Such  bodies  make  good  dehydrating,  i.  e., 
drying,  agents. 

Examples  are:  Anhydrous  cupric  sulphate  and  anhydrous 
potassium  carbonate." 

If  dehydrating  agents  absorb  so  much  water  from  the 
air  that  they  dissolve  in  the  water,  they  are  said  to 
deliquesce. 

An  example  is  anhydrous  calcium  chloride. 

In  any  case,  if  a  substance  takes  up  water  when  ex- 
posed to  moist  air,  it  is  said  to  be  hygroscopic. 

Many  substances  are  drying  agents,  not  because  they 
tend  to  take  up  water  of  crystallization,  but  because  they 
combine  with  water  to  form  other  compounds.  Quicklime, 
which  is  calcium  oxide,  is  an  example.  When  this  sub- 
stance is  slaked,  by  addition  of  water,  it  becomes  calcium 
hydroxide. 

f 
54.  Exercises.  I 

1.  How  "many  grams  of  water  are  formed  by  the  combustion 
of  10  grams  of  hydrogen  in  air  ? 

2.  What  evidence  is  there  that  the  hydrogen  of  water  is 
more  divisible  than  the  oxygen  ? 

3.  How  could  you  determine  approximately  how  much  water 
is  contained  in  a  potato  ? 


EXERCISES.  57 

4.  5  grams  of  crystalline  barium  chloride  lost,  when  heated 
at  120°  C.,  0.65  grams.     What  per  cent  of  water  did  it  have  ? 

5.  Calculate  the  per  cent  of  water  of   crystallization   in   a 
sample  of  potash  alum,  47.4  grams  of  which  lost. 21. 6  grams 
of  water. 

6.  How   many  grams  of  water  can  be  decomposed  by  10 
grams   of    sodium?     How  much    sodium    hydroxide    will   be 
formed  ? 

7.  How  many  grams  of  potassium  are  required  to  give  with 
water  50  grams  of  potassium  hydroxide?     How  many  grams 
of  hydrogen  will  be  liberated  at  the  same  time  ?     How  many 
liters  when  1  liter  weighs  0.09  grams? 

8.  How   many  grams  of   water   are  formed  by  burning  10 
liters   of  hydrogen  when  1   liter  of  the  latter  weighs   0.085 
grams  ? 


CHAPTER  IV. 
SOLUTION. 

55.  The  Character  of  Solution.  —  Solution,  or  dis- 
solving, takes  place  when  substances  are  mixed  in  such 
a  way  that  the  matter  of  each  is  distributed  uniformly 
through  that  of  the  others.  The  resulting  homogeneous 
mixture  is  called  a  solution.  Thus  broadly  denned,  the 
term  solution  includes  phenomena  called  by  many  dif- 
ferent names,  but  we  shall  restrict  it  to  the  absorption 
of  a  gas,  liquid,  or  solid,  within  the  portion  of  space 
occupied  by  some  liquid.  The  liquid  is  called  the 
solvent.  Examples  of  common  solvents  are :  Water, 
alcohol,  and  ether. 

If  the  solvent  is  colorless,  and  the  dissolved  substance  has  a 
definite  color,  the  solution  will  usually  be  colored  ;  if  the  dis- 
solved substance  is  white,  or  colorless,. the  solution  will  be 
colorless  ;  but  in  any  case  the  solution  will  be  clear.  Insoluble 
substances  often  remain  mechanically  suspended  in  a  liquid  for 
some  time.  Their  presence  is  shown  by  the  turbid  appearance 
of  the,  liquid. 

Dilute  solutions  have  practically  the  same  volume  as 
that  of  the  solvent,  but  when  the  amount  of  dissolved 
substance  becomes  large,  the  volume  of  the  solution  is 
increased. 

68 


TEMPERATURE   CHANGES.  59 

An  illustration  of  the  first  statement  is  found  in  the  familiar 
experiment  in  which  a  considerable  quantity  of  powdered 
sugar  is  added,  little  by  little,  to  a  vessel  entirely  full  of  water 
without  causing  an  overflow,  while  a  much  smaller  amount  of 
an  insoluble  substance,  e.  g.,  sand,  causes  some  of  the  water  to 
be  displaced. 

56.  Boiling  Point  and  Freezing  Point  of  a  Solution. 

Substances  in  solution  raise  the  boiling  point,  and  lower 

the  freezing  point  of  the  solvent.  Thus,  water  con- 
taining salt  or  sugar  boils  above  100°  C.  at  760  mm. 
pressure,  and  freezes  below  0°  C.  In  dilute  solutions 
the  rise  of  the  boiling  point  and  the  lowering  of  the 
freezing  point  are  proportional  to  the  amount*  of  dissolved 
substance  in  a  given  volume.  The  specific  gravity  of 
solutions  of  solids  is  greater  than  that  of  the  solvent. 
Thus,  sea  water  has  a  specific  gravity  of  1.026. 

57.  Temperature  Changes  during  Solution.  — When 
a  gas  dissolves  in  a  liquid  there  is,  as  a  rule,  an  evolu- 
tion of  heat  and  a  consequent  rise  of  temperature,  but 
the  solution  of  a  solid  is  usually  attended  by  an  absorp- 
tion of  heat  and    a   reduction  of  temperature.     Some 
solids,    however,  dissolve    in  water  with    evolution   of 
heat.     Examples  are :  Anhydrous  calcium  chloride  and 
anhydrous    sodium    carbonate.     These  apparent  excep- 
tions are  usually  substances  that  have  been  deprived  of 
water   of   crystallization,    and  take   it  up  again  when 
brought  into  contact  with  water.     Because  of  the  heat 


60  SOLUTION. 

evolution  due  to  the  union  of  these  substances  with 
their  crystal-water,  the  heat  absorption  due  to  the 
solution  of  the  crystallized  substances  is  not  perceptible. 
When  the  crystals  of  such  substances  are  dissolved  in 
water  there  is  usually  a  reduction  of  temperature. 


58.  Solubility.  —  By  the  solubility  of  a  substance 
we  mean  the  maximum  amount  of  the  substance  that 
can  be  taken  up  by  a  given  quantity  of  the  solvent 
under  the  given  conditions. 


The  amounts  of  two  substances  which  will  dissolve  in  a 
given  weight  of  a  solvent  are  very  unequal,  as  are,  also,  the 
quantities  of  two  solvents  which  are  required  to  absorb  a  given 
weight  of  the  same  substance.  Thus,  sugar  and  salt  are  very 
soluble  in  water,  but  practically  insoluble  in  ether.  Even  in 
water,  however,  their  solubilities  are  very  different,  sugar  being 
much  more  soluble  than  salt.  Similar  differences  exist  in  the 
case  of  gases,  hydrogen,  for  example,  being  only  about  half  as 
soluble  as  oxygen  in  water  of  the  ordinary  temperature. 


59.  Effect  of  Temperature  on  Solution.  --  The  solu- 
bility of  a  substance  depends  not  only  upon  the  solvent 
used,  but  also  upon  the  temperature.  As  a  rule,  solids 
are  more  soluble  in  hot  than  in  cold  liquids,  while  the 
reverse  is  true  of  gases.  The  following  table  shows  the 
effect  of  temperature  upon  the  solubility  of  several 
solids, 


SOLUBLE  AND  INSOLUBLE  SUBSTANCES. 


61 


SUBSTANCE. 

GRAMS  SOLUBLE  IN  100  GRAMS  WATER. 

AtO°C. 

At  20°. 

At  100°. 

Potassium  nitrate. 

13.3 

31.7 

246. 

Sodium  chloride. 

35.    " 

36. 

39.7 

Potassium  chlorate. 

7.2 

59.5 

Cupric  sulphate  (cryst.). 

42.3 

203.3 

To  illustrate  the  decrease  of  solubility  of  gases  with 
rise  of  temperature  we  may  take  the  case  of  oxygen, 
4.1  c.c.  of  which  can  dissolve  in  100  c.c.  of  water  at 
0°  C.,  2.9  c.c.  at  15°,  and  none  at  100°. 

60.  Soluble  and  Insoluble  Substances.  —  A  solid 
requiring  less  than  100  times  its  weight  for  com- 
plete solution  may  be  considered  readily  soluble ; 
one  needing  between  100  and  1,000  times  its  weight, 
difficultly  soluble ;  while  one  which  requires  more  than 
1,000  parts  of  the  solvent  may  be  called  insoluble. 

There  is,  however,  a  great  difference  in  the  solubilities  of 
so-called  "  insoluble  "  substances.  Thus,  strontium  sulphate 
requires  about  8,000  parts  of  water  for  solution,  and  barium 
sulphate,  about  400,000  parts. 

We  call  a  substance  insoluble,  then,  only  relatively  to 
other  substances,  and  not  absolutely,  for  there  is  prob- 


62  SOLUTION. 

ably  no  substance  of  which  a  small  amount  will  not  dissolve, 
if  the  quantity  of  the  solvent  is  very  large. 

61.  Saturated  Solution.  —  When  a  liquid  has  in  solu- 
tion all  that  it  can  hold  of  a  substance  under  certain 
conditions,  it  is  said  to  be  saturated  with  respect  to  that 
substance   under  the    specified    conditions.     Considering 
now  only  the  solution  of  solids  in  liquids,  we  may  know 
that  a  solution  is  saturated  when  a  slight  lowering  of 
its  temperature  or  the  removal  of  a  small  amount  of  the 
solvent,  e,  g.,  by  evaporation,  causes  precipitation  of  some 
of  the  dissolved  substance. 

There  are  two  methods  in  common  use  for  the  production, 
at  ordinary  temperatures,  of  a  saturated  solution  of  a  solid. 
In  the  first  of  these  the  solvent  is  allowed  to  remain  for  some 
time  in  contact  with  an  excess  of  the  solid,  and  the  mixture  is 
shaken  or  stirred  to  hasten  solution. 

In  the  second  meth'od,  the  solvent  is  heated  above  the  ordi- 
nary temperature  with  enough  solid  to  produce  saturation  at 
the  higher  temperature,  and  the  solution  is  then  cooled  to  the 
ordinary  temperature.  In  this  way,  the  excess  of  solid  is 
deposited. 

62.  Supersaturated    Solutions. -- Many    solutions, 
however,  although  saturated  at  temperatures  above  the 
ordinary,  will  not  deposit  their  excess  of  dissolved  solid 
when  the  temperature  is  lowered.     Such  solutions  are 
said  to  be  supersaturated.     A  solution  usually  remains 
supersaturated  only  while  undisturbed.     If  the  contain- 
ing vessel  is  jarred,  or  if  small  particles,  e.  #.,  of  dust, 


P&ECIPiTATION  AND    CttYSTALLlZAflOfr.  63 

are  introduced,  precipitation  often  results.  The  most 
certain  way,  however,  of  disturbing  a  supersaturated  so- 
lution is  to  add  a  crystal  of  the  dissolved  substance.  A 
rapid  separation  of  the  excess  of  the  latter  is  the  result. 

63.  Precipitation  and  Crystallization.  —  As  stated  in 
the  preceding  section,  a  solid  may  be  made  to  separate 
from  its  solution,  if  the  latter  is  brought  to  the  point  of 
saturation.  A  dilute  solution  must,  therefore,  be  con- 
centrated if  separation  is  to  take  place.  If  a- solid  sep- 
arates from  solution  rather  slowly,  it  will  frequently  be 
found  to  consist  of  regular  masses  called  crystals.  The 
more  sloivly  crystallization  takes  place,  the  larger  and 
more  perfect  will  the  crystals  be.  But  often  a  solution  is 
brought  to  saturation  suddenly,  as  is  the  case  when  the 
temperature  of  an  almost  saturated  solution  is  rapidly 
lowered,  or  when  another  solvent  is  added,  or  when  a 
new  substance  is  formed  which  is  not  very  soluble  in 
the  solvent.  In  such  cases,  the  substance  that  separates 
from  solution  will  consist  of  very  small  crystals,  or  it 
may  even  be  in  an  amorphous,  i.  e.,  non-crystalline,  form. 
In  either  case  it  is  called  a  precipitate. 


Thus,  when  silver  nitrate  and  sodium  chloride  solutions  are 
mixed,  there  is  produced  a  white,  amorphous  precipitate  of 
silver  chloride,  the  sodium  nitrate  formed  at  the  same  time 
remaining  in  solution.  Relatively  to  silver  nitrate  solution, 
therefore,  sodium  chloride  solution  is  a  precipitant,  since  its 
addition  produces  a  precipitate. 


64  SOLUTION. 

64.  Decantation  and  Filtration.  —  We  may  separate 
a  precipitate  from  the  solution  in  which  it  is  suspended 
either  by  allowing  it  to  settle  and  then  decanting,  i.  e., 
pouring  off,  the  clear  solution,  or  by  filtering  the  mixture 
of  liquid  and  solid.     For  the  latter  purpose  we  use  a 
filter  paper,  consisting  of  cellulose,  which  permits  liquids 
and  dissolved  solids  to  pass  through  its  pores,  but  usu- 
ally holds  back  suspended  solids.     What  passes  through 
the  filter  is  called  the  filtrate. 

65.  Crystallization  from  Fusion.  —  A  solid  may  sep- 
arate in  crystalline-  form  not  only  from  solution,  but  also 
by  the  solidification  of  a  liquid,  i.  e.,  from  fusion.     Thus, 
water  freezes,  and  molten  sulphur  solidifies,  in  crystal- 
line form. 

Just  as  there  is  a  supersaturated  condition  of  some  so- 
lutions, owing  to  a  tardy  separation  of  dissolved  solids, 
so  there  is  a  super/used  condition  of  some  liquids, 
because  of  their  slow  assumption  of  the  solid  form  even 
at  temperatures  below  their  true  freezing  points.  Crys- 
tallization is  effected  in  the  same  way  in  both  cases,  viz., 
by  "  inoculation  "  with  a  crystal  of  the  solid. 

66.  Effervescence.  —  Gases,  like  solids  and  liquids, 
separate  from  solution  if  formed  in  a  solvent  unable  to 
hold  them.     Because  of  their  low  specific  gravity,  how- 
ever, gases  rise  to   the   top   of  the  solution,  and   thus 
escape  into  the  air.     A  liquid  evolving  a  gas  is  said  to 
effervesce. 


EXERCISES.  65 

The  action  of  zinc  on  dilute  sulphuric  acid,  for  example, 
causes  effervescence  of  the  dilute  acid,  owing  to  the  escape  of 
hydrogen.  Similarly,  "soda  water"  effervesces,  because  of 
liberation  of  carbon  dioxide  gas. 

67.  Exercises. 

1.  Why  do  subterranean  waters  contain  more  gaseous  sub- 
stances in  solution  than  surface  waters  ? 

2.  Why  is  it  that  insoluble  substances,  e.  </.,  sand,  "  burnt" 
alum,  etc.,  have  no  taste  ? 

3.  Suggest  a  method  of  separating  a  mixture  of  potassium 
nitrate  and  sodium  chloride  so  as  to  recover  almost  all  of  the 
nitrate.     See  the  table  of  solubilities  in  §  59. 

4.  Why  does  "  spattering  "  take  place  in  the  evaporation  to 
dry  ness  of  a  solution  of  common  salt,  etc.  ? 

5.  How  could  you  separate  a  mixture  of  white  sand  and 
common  salt  so  as  to  recover  all  of  each  in  a  dry  condition  ? 

6.  Calculate  the  parts  per  cent  of   potassium  dichromate 
present  in  a  solution  containing  120  grams  of  water  and  15 
grams  of  the  dichromate. 


CHAPTER   V. 

FUNDAMENTAL  LAWS,  COMBINING   NUMBERS  AND 
NOMENCLATURE. 

68.  Persistence  of  Mass.  —  The  fundamental  fact 
regarding  every  chemical  change  is  that  it  results  in  the 
formation  of  at  least  one  new  substance.  If  the  new 
substance  is  an  element,  it  can  have  been  formed  only 
by  the  decomposition  of  some  previously  existing  com- 
pound; if  a  compound,  only  by  the  union  of  certain 
elements.  Thus  considered,  every  chemical  reaction  is 
merely  a  change  in  the  relations  between  elementary  sub- 
stances. 

Most  of  the  reactions  already  studied  may,  in  fact,  be 
classified  under  one  of  three  heads :  — 

1 .  Elements  in  combination  became  separated  by  a  change 
in  conditions;  as  when  mercuric  oxide  was  decomposed  by 
heat  into  mercury  and  oxygen. 

2.  Element*  (or  an  element  and  a  compound)  existing 
apart  united  to  form  a  (or  another)  compound;  as  when 
hydrogen  and  oxygen  combined  to  form  water. 

3.  A  free  element  took  the  place  of  one  of  the  elements 
of  a  compound,  the   latter    element   being    thereby    freed 
from  combination.     An  illustration  is  the  case  of  zinc  and 
dilute    sulphuric   acid.     Here    hydrogen    is    set   free   from 
the  acid,  and  zinc   enters   into    combination   in   the   place 
of  the  hydrogen. 


CONSTANT  PROPORTIONS.  67 

Since,  therefore,  chemical  changes  are  only  re-arrange- 
ments of  elements  already  existing,  the  sum  of  masses 
of  the  reacting  substances  must  always  be  equal  to  the  sum 
of  the  masses  of  the  products.  This  is  the  law  of  "  Per- 
sistence of  Mass,"  or  "  Conservation  of  Matter."  It 
has  proved  to  be  true  in  every  case  that  has  been 
examined, 

69.  Constant  Proportions.  —  There  is  another  law  of 
chemical  action  that  is  related  closely  to  the  law  of  Per- 
sistence of  Mass;  it  is  the  law  of  "  Constant  Propor- 
tions by  Weight."  This  law,  like  the  first,  is  a  general 
statement  of  facts  learned  from  many  experiments.  It 
may  be  stated  thus  :  "The  relation  between  the  masses  of 
reacting  substances,  and  between  them  and  the  masses  of 
the  products,  is  definite  and  constant." 

All  the  chemical  reactions  previously  studied  illustrate  this 
law.  Thus,  122.5  grams  of  potassium  chlorate  always  give, 
when  completely  decomposed,  74.5  grams  of  potassium  chloride 
and  48  grams  of  oxygen.  Another,  illustration  is  the  case  of 
sodium  and  water,  which  always  react  in  the  proportion  of  23 
grams  of  sodium  to  18  of  water,  giving  40  grams  of  sodium 
hydroxide  and  1  of  hydrogen.  JVb  accurate  experiment  has  ever 
shown  that  23  grams  of  sodium  liberate  any  other  quantity  of 
hydrogen  than  approximately  1  gram. 

Another  way  in  which  this  law  may  be  stated  is :  A 
given  chemical  compound,  no  matter  what  its  source,  will 
always  be  found  to  be  composed  of  the  same  elements  united 
in  the  same  proportions. 


68  FUNDAMENTAL  LAWS. 

To  illustrate  :  All  determinations  of  the  composition  of 
water  show  that  the  relation  of  about  11.11%  hydrogen  to 
88.89%  oxygen  is  always  preserved.  So,  too,  mercuric  oxide  is 
always  composed  of  mercury,  92.59%,  and  oxygen,  7.41%  ;  and 
sulphur  dioxide  of  sulphur,  50%,  and  oxygen,  50%. 

70.  Symbols  and  Formulas. — In  all  the  reactions 
hitherto  studied  we  have  written  out  in  full  the  names 
of  the  substances  taken  and  obtained,  together  with  the 
proportion    by  weight  of    each.     As   a  rule,  however, 
chemists  use   symbolic  expressions  in  place  of   the  full 
names  of  substances.     When  this  method  is  once  un- 
derstood it  will  be  seen  to  have  many  advantages.     As 
has  already  been  stated  in  the  introductory  chapter,  the 
elements  may  be  represented  by  symbols. 

A  symbol  is  simply  the  initial  letter,  or  the  initial  fol- 
lowed by  another  characteristic  letter,  of  the  name  of  the 
element.  We  indicate  the  composition  of  a  compound 
substance  by  writing  the  symbols  of  the  elements  which 
make  up  the  compound  side  by  side,  without  an  inter- 
vening sign  (-f-,  — , >  ,  or  =).  The  resulting  ex- 
pression is  called  a  formula. 

Thus,  mercuric  oxide,  a  compound  of  mercury  and  oxygen, 
is  represented  by  the  formula  HgO.  Similarly,  HC1  represents 
a  compound  of  hydrogen  and  chlorine,  i.  e.,  hydrochloric  acid. 

71.  Symbolic    Equations. — A    symbolic     equation 
(called,  simply,  "  an  equation ")  is  formed  by  writing, 
instead  of  the  names  of  the  reacting  substances,  their 


SYMBOLIC  EQUATIONS.  69 

symbols  and  formulas.  Thus,  the  fact  that  sodium  acts 
upon  water  to  produce  sodium  hydroxide  and  hydrogen 
is  shown  by  the  expression, 

Na  +  HOH  (or  H2O) »  NaOH  +  H. 

In  this  equation  Na  is  the  symbol  of  sodium ;  H,  that  of 
hydrogen ;  and  O,  that  of  oxygen.  The  formula  HOH, 
or  H2O,  for  water,  shows  that  water  is  composed  of 
hydrogen  and  oxygen,  while  the  formula  NaOH,  for 
sodium  hydroxide,  shows  that  sodium  hydroxide  is  a 
compound  of  sodium,  oxygen,  and  hydrogen.  The  sign 

»  (or  =)  is  best  read  "give,"  or  "produce."     The 

sign  -f-  is  read  "and."     The  symbols  and  formulas  of 

the  equation  are  read  in  the  direction  of  the  arrow » ; 

those  preceding  the  arroAV  are  called  the  factors,  those 
succeeding  it  the  products  of  the  reaction.  When  the 
sign  =  is  used,  the  symbols  and  formulas  to  the  left  of 
it  are  the  factors,  and  those  to  the  right  the  products. 


Among  the  advantages  of  the  use  of  symbols  are  the  fol- 
lowing :  — 

1.  Symbolic  expressions  for  compounds  enable  us  to  see,  at 
once,  what  elements  make  up  the  compounds.     Thus,  while  the 
common  names  of  salt,  water,  galena,  and  caustic  soda  give  us 
no  idea  of  the  composition  of  these  bodies,- the  corresponding 
formulas,  NaCl,  II2O,  PbS,  and  NaOH,  do. 

2.  Symbolic  equations,  if  correctly  written,  enable  us  to  de- 
termine what  changes  in  the  relations  of  elements  have  resulted 
from  a  given  case  of  chemical  action.     Thus  it  requires  only  a 
glance  at  the  symbolic  equation  for  the  action  of  sodium  upon 


70  FUNDAMENTAL  LAWS. 

water  to  determine  that  in  this  reaction  sodium  displaced  part 
of  the  hydrogen  of  water  to  form  sodium  hydroxide. 

72.  Equations  the  Result  of  Experiment. — Equa- 
tions mean  nothing  unless  they  are  the  result  of  experiment. 
If  the  student  cannot  himself  prove  what  products  are 
formed  in  a  reaction,  he  must  depend  upon  some  trust- 
worthy source,   e.  g.,    the    teacher   or   a   text-book,   for 
the  necessary  information.     The  student  should  learn 
the  important  equations  given  from  time  to  time  in  the 
text,  hot  by  rote,  but  with  appreciation  of  their  mean- 
ing.    He  will  soon  find  that  the  number  of  equations 
that  must  actually  be  memorized  is  small,  for,  with  a 
few  typical   equations    as   a   basis,  a  large   number   of 
analogous  ones  can  readily  be  acquired. 

73.  Quantitative  Meaning  of  Symbols  and  Equa- 
tions. —  The  equation 

Na  +  H2O »  NaOH  +  H 

means  to  the  chemist  much  more  than  has  been  stated 
in  §  10,  for  it  indicates  the  proportions  by  weight  of  the 
reacting  substances  and  of  the  products.  We  can  give 
this  added  meaning  to  the  above  and  to  every  equation 
by  letting  the  symbol  of  each  element  represent  not 
only  the  element  in  general,  but  also  a  definite  mass  of 
it.  The  formula  of  a  compound  will  then  show  not 
only  what  elements  are  contained  in  the  compound,  but, 
in  addition,  the  proportions  by  iveight  in  which  they  are 
united.  Finally,  if  we  assume  that  symbols  and  forniu- 


SYMBOLS  AND  FORMULAS.  71 

las  stand  for  definite  masses  of  elements  and  of  com- 
pounds, respectively,  then  the  equation  for  every  reaction 
will  represent  the  proportion  by  weight  of  every  substance 
entering  into  the  reaction. 

Thus,  the  equation 

Ka  +  H2O »  NaOH  +  H 

means  to  the  chemist  that  23  parts  by  weight  of  sodium  react 
with  18  parts  of  water  to  produce  40  parts  of  sodium  hydroxide 
and  1  part  of  hydrogen.  Of  course  the  proportions  will  be 
true,  whatever  units  are  used,  i.  e.,  with  pounds  and  tons  as 
well  as  with  milligrams  and  grams,  but  for  our  present  purpose 
we  will  let  symbols  and  formulas  represent  grams.  Na  means, 
therefore,  23  grams  of  sodium,  and  H2O,  18  grams  of  water. 
Since  water  is  one-ninth  hydrogen  and  eight-ninths  oxygen,  18 
grams  of  water  must  contain  2  grams  of  hydrogen  and  16  of 
oxygen.  O,  therefore,  means  16  grams  of  oxygen. 

74.  How  to  Represent  Multiples  of  Symbols  and 
Formulas.  —  In  the  formula  H2O,  H2  means  twice  the 
quantity  of  hydrogen  represented  by  H,  for  in  all  chemi- 
cal formulas  a  small  figure  written  after  and  slightly 
below  a  symbol  multiplies  the  quantity  represented  by 
the  symbol  immediately  preceding. 

The  formula  NaOH  means  40  grams  of  sodium 
hydroxide.  Of  the  40  grams,  23  are  sodium,  16  are 
oxygen,  and  1  is  hydrogen.  The  equation  thus  ac- 
counts for  every  gram  of  material  taken:  — • 

Na  +  H2O »  NaOH  +  H. 

23  +  (2  +16)          (23  +16  +1)  +1 


72  FUNDAMENTAL   LAWS. 

If  we  multiply  the  quantities  taken  in  the  above  equation 
by  2,  the  relative  amounts  are  not  altered.  Thus  the  equatior 

2  Ka-f  2  H2O »  2  KaOH  +  H2  represents  all  the  facts  as 

well,  at  least,  as  the  simpler  equation. 

A  figure  written  before  a  formula  multiplies  the 
quantity  of  the  substance  represented  by  the  formula, 
just  as  the  small  figure  written  after  a  symbol  multiplies 
the  quantity  indicated  by  the  symbol. 

Therefore  2  ]STa  means  46  (=  23  X  2)  grams  of  sodium  ;  2 
H2O  means  36  (=2  [2  +  16]  )  grams  of  water  ;  2  XaOH  means 
80  (=2  [23  -j-  16-j-l]  )  grams  of  sodium  hydroxide  ;  and  H2 
means  2  grams  of  hydrogen. 

It  often  happens  that  a  formula  contains  a  group  of 
elements  repeated  two  or  more  times.  Thus,  the  for- 
mula of  calcium  hydroxide  is  Ca(OH)2  ;  it  may  also 

OH 
be    written    CaO2H2 ;  or,  better  still,  Ca^,j.     In    the 

first  of  these  formulas,  viz.,  Ca(OH)2,  the  small  figure 
written  after  the  parenthesis  multiplies  what  is  in  the  par- 
enthesis immediately  preceding,  just  as  if  the  symbols  in 
the  parenthesis  were  together  one  symbol.  Symbols 
which  are  grouped  together  in  this  way  are  called 
radicals. 

Other  examples  of  formulas  containing  radicals  are  :  Cu- 
(XO3)2  for  cupric  nitrate,  and  Fe2(SO4)3  for  ferric  sulphate. 
These  might  be  written  Cu!N"2O6  and  Fe2  83O12,  respectively, 
but  the  first  formulas  are  preferable  because  they  enable  us  to 


COMBINING  PROPORTIONS. 


73 


see  that  the  compounds  are  derived  /rom  nitric  acid  and  sul- 
phuric acid,  HNO3  and  H2SO4,  respectively,  while  the  second 
formulas  do  not. 

The  water  of  crystallization  present  in  many  com- 
pounds is  represented  by  the  formula  of  water  (taken 
the  necessary  number  of  times)  after  the  formula  of 
the  compound.  Thus,  while  CuSO4  stands  for  anhy^ 
drous  cupric  sulphate,  blue  vitriol  has  the  formula  Cu- 
SO4.  5  H2O  ;  crystallized  zinc  sulphate  is  ZnSO4.  7  H2O: 
and  crystallized  sodium  sulphate  (Glauber's  salt)  is 
Na2S04.  10  H20. 

75.  Combining  Proportions.  —  The  following  list 
gives  the  names  of  some  of  the  more  important  elements, 
and  the  proportions  by  weight,  or  a  sub-multiple  of  them, 
in  ivhich  these  elements  unite  when  they  form  compounds. 


s  ^ 

J 

£ 

fc  2 

ELEMENT. 

ELEMENT. 

i| 

CO 

i  | 

N 

co 

H 

Aluminum. 

Al. 

27. 

Manganese. 

Mn. 

55. 

Barium. 

Ba. 

137. 

Mercury. 

Hg. 

200. 

Bromine. 

Br. 

80. 

Nitrogen. 

No 

14. 

Calcium. 

Ca. 

40. 

Oxygen. 

0. 

16. 

Carbon. 

C. 

12. 

Phosphorus. 

P. 

31. 

Chlorine. 

Cl. 

35.5 

Potassium. 

K. 

39. 

Copper. 

Cu. 

63. 

Silicon. 

Si. 

28. 

Fluorine. 

Fl. 

19. 

Silver. 

Ag. 

108. 

Hydrogen. 

H. 

1. 

Sodium. 

Na. 

23. 

Iodine. 

I. 

127. 

Strontium. 

Sr. 

87. 

Iron. 

Fe. 

56. 

Sulphur. 

S. 

32. 

Lead. 

Pb. 

207. 

Tin. 

Sn. 

118. 

Magnesium. 

Mg. 

24. 

Zinc. 

Zu. 

65.    A 

74  FUNDAMENTAL   LAWS. 

76.  The    Use   of   Combining    Proportions.  —  The 

derivation  of  the  combining  proportions  given  above  is 
of  great  importance  to  Chemistry ;  however,  it  depends 
upon  facts  which  we  are  not  ready  to  consider  at  this 
time.  For  the  present,  the  combining  proportions  given 
in  the  preceding  table  will  enable  us  to  use  equations, 
formulas,  and  symbols  quantitatively. 

Thus,  if  the  formula  MgO  is  given  to  the  compound  which 
magnesium  forms  when  it  burns  in  oxygen,  a  reference  to  the 
table  will  show  that  in  this  compound  magnesium  and  oxygen 
are  united  in  the  proportion  of  24  grams  of  magnesium  to  16 
grams  of  ox}-gen.  From  24  grams  of  magnesium,  therefore, 
we  can  get  40  grams  of  magnesium  oxide.  Again,  if  A12O3 
represents  correctly  the  composition  of  aluminum  oxide,  the 
student  can  see  readily  that,  in  this  compound,  aluminum  and 
oxygen  are  combined  in  the  proportion  of  54  grams  of  alumi- 
num to  48  of  oxygen. 

The    equation    Zn  +  H2SO4 »  ZnSO4  +  H2  thus 

means,  as  the  result  of  a  simple  calculation,  that  65 
grams  of  zinc  react  with  98  grams  of  sulphuric  acid  to 
give  161  grams  of  zinc  sulphate  and  2  grams  of  hydro- 
gen. 

77.  How  a  Compound  of  Two  Elements  is  Named. 

—  A  compound  of  two  elements  lias  the  name  of  each  ap- 
pearing in  its  name,  but  the  last  syllable  of  the.  name  of 
one  of  the  elements  is  changed  to  ide. 

Thus  sodium  chloride  is  a  compound  of  sodium  and  chlorine  ; 
magnesium  oxide,  one  of  magnesium  and  oxygen  (here  the  y 


HOW   TO  DISTINGUISH  COMPOUNDS.  75 

before   the  ide  is   omitted)  ;  calcium  carbide,  one  of  calcium 
and  carbon  ;  and  lead  sulphide,  one  of  lead  and  sulphur. 

As  to  which  element  shall  have  the  ending  ide,  the 
rule  is  as  follows :  If  one  of  the  elements  is  an  un- 
doubted metal,  it  retains  its  full  name,  as  in  the  cases 
given  above,  and  the  ending  of  the  non-metal  is  changed 
to  ide.  If  neither  is  a  metal,  as  in  the  case  of  a  com- 
pound of  sulphur  with  oxygen,  the  name  of  the  one 
having  the  smaller  resemblance  to  a  metal  is  changed  to 
ide.  Thus,  a  compound  of  sulphur  with  oxygen  is 
called  sulphur  oxide,  but  one  of  sulphur  and  hydrogen 
is  called  hydrogen  sulphide.  (N.  B.  Hydrogen  has 
certain  metallic  properties.)  If  both  elements  were 
metals,  the  same  rule  would  hold,  and  the  one  which 
possessed  the  less  characteristically  metallic  properties 
would  have  its  ending  changed  to  ide. 

78.  How  to  Distinguish  between  Compounds  of  the 
Same  Two  Elements.  —  If  there  is  more  than  one  com- 
pound of  the  same  two  elements,  the,  ending  of  the  less 
metallic  element  is  changed  in  each  case,  as  before,  to 
ide,  but  the  compounds  are  distinguished  from  one 
another  in  one  of  two  ways.  These  are,  — 

(1)  By  changing  the  final  syllable  of  the  more  metallic  (i.  e.9 
electro-positive)  element ;  or, 

(2)  By  placing  a  prefix  before  the  name  of  the  less  metallic 
(i.  e.j  more  electro-uegative)  element. 


76  FUNDAMENTAL  LAWS. 

By  the  first  method  the  ending  of  the  metallic  ele- 
ment is  changed  to  ous  or  ic,  as,  for  example,  in  the 
names  of  the  two  compounds  of  mercury  and  chlorine, 
which  are  called  mercurowa  chloride  and  mercuric  chlor- 
ide, respectively.  The  ending  ous  is  given  to  that  one 
of  the  compounds  which  contains  the  larger  proportion 
of  the  metallic  element  (here  mercury);  the  ending  ic, 
on  the  other  hand,  to  the  one  containing  the  smaller 
proportion  of  the  metallic  element. 


Ous  means  in  chemical  phraseology,  as  in  ordinary  language, 
"full  of"  or  "containing  much  of."  The  appropriateness  of 
the  names  for  the  two  compounds  of  mercury  and  chlorine  will 
be  seen  by  a  comparison  of  the  proportionate  amounts  of  the 
elements  in  the  two  compounds.  Thus,  mercurous  chloride, 
HgCl,  has  200  parts  by  weight  of  mercury  to  35.5  of  chlorine, 
while  mercuric  chloride,  HgCl2,  has  200  of  mercury  to  71  of 
chlorine.  The  larger  proportion  of  mercury  is,  evidently,  in 
the  mercurous  compound. 


In  the  case  of  some  elements,  e.  g.,  copper  and  iron, 
the  Latin  names  are  used,  and  the  ending  is  applied  to 
these  rather  than  to  the  English  names.  Thus,  the  two 
common  copper  oxides  are,  — 


(a)  Cuprous  oxide,  containing  copper,  63  parts,  to  oxygen, 
8  parts  ; 

(6)  Cuprtc  oxide,  containing  copper,  63  parts,  to  oxygen,  16 
parts. 


EXERCISES.  77 

Similarly,  iron  compounds  are  distinguished  by  the 
names,  — 

(a)  Ferrous  oxide,  chloride,  etc.,  for  the  compound  contain- 
ing the  larger  proportion  of  iron  ;  and 

(6)  Feme  oxide,  chloride,  etc.,  for  the  one  containing  the 
smaller  proportion  of  iron. 

A  second  way  of  distinguishing  between  two  (or 
more)  compounds  of  the  same  two  elements  is  to  apply 
a  numerical  prefix  to  the  less  metallic  element,  and  to 
leave  the  metallic  element  unchanged.  Thus  the  names 
carbon  monoxide  and  carbon  o&oxide  distinguish  the 
two  compounds  of  carbon  and  oxygen  from  each  other. 

The  prefix  mon  means  "  one  "  or  "  first,"  referring  to  the 
formula  CO,  while  di  means  "two"  or  "second,"  referring 
to  the  formula  CO2. 

In  a  similar  way  we  distinguish  sulphur  dioxide,  SO2, 
from  sulphur  trioxide,  SO3. 

79.  Exercises. 

1.  Calculate  the  percentage  composition  of  sulphuric  acid, 
sodium  hydroxide,  manganese  dioxide  (MnO2),  and  potassium 
chlorate  (KC1O3). 

2.  What   relative   quantities  of   the  substances   taken   and 
produced  are  indicated  in  the  equations, 

2  KC1O3  — »  2  KC1  -f  3  O2,  and 
2  K  -f-  2  H2O >  2  KOH  +  H2  ? 


78  FUNDAMENTAL  LAWS. 

3.  What  quantity  of  hydrogen  could  be  obtained  from  the 
action  of  50  grams  of  potassium  upon  an  excess  of  water  ? 

4.  Calculate  the  parts  per  cent  of  water  of  crystallization  in 
blue  vitriol,  in  Glauber's  salt,  and  in  gypsum,  CaSO4.  2  H2O. 

5.  What  weight  of  anhydrous  zinc  sulphate  is  contained  in 
75  grams  of  the  crystallized  form,  ZnSO4.  7  H2O  ? 

6.  How    many  grams    of  phosphorus   must   be   burned   to 
produce  60  grams  of  phosphorus  pentoxide,  P2O5  ? 

7.  Name   the   compounds  having   the   formulas   CO,   CS2, 
ZnCl2,  Hgl,  and  BaC8. 


CHAPTER   VI. 
CHLORINE. 

80.  Existence.  — Chlorine  is  a  heavy,  greenish-yellow 
gas,  of  irritating  odor  and  poisonous  properties.     It  was 
discovered  by    the  Swedish  chemist  Scheele  in  1774, 
but  was  not  generally  considered  to  be  an  element  until 
1809.    . 

Because  of  its  great  reactivity,  i.  e.,  its  tendency  to 
act  chemically  with  other  substances,  chlorine  is  not 
found  in  nature  free,  but  always  in  combination  with 
other  elements.  Its  most  abundant  compounds  are 
sodium,  potassium,  and  magnesium  chlorides  and  hy- 
drochloric acid,  which  is  hydrogen  chloride.  Sodium 
chloride  is  common  salt. 

81.  Common  Method  of  Preparation.  —  Chlorine  is 
usually    prepared    by    the    action    of    reagent    hydro- 
chloric acid  upon  manganese  dioxide.     The  apparatus 
is  shown  in  Fig.  19. 

A  flask  containing  manganese  dioxide  (MnO2)  in  small 
lumps  is  provided  with  a  thistle  tube  and  a  delivery  tube,  and 
is  supported  so  that  it  may  be  warmed  in  a  water  bath.  Con- 
centrated hydrochloric  acid  is  added  through  the  thistle  tube, 
and  the  evolution  of  chlorine  begins.  The  gas  is  allowed  to 
pass  through  a  wash  bottle  containing  a  little  water,  a  drying 

79 


80 


CHLOEINL'. 


bottle  one-third  full  of  concentrated  sulphuric  acid  (a  U-tube 
of  calcium  chloride  may  be  used  instead),  and  then  into  a  col- 
lecting bottle,  the  air  of  which  is  to  be  displaced  by  chlorine. 
From  the  collecting  bottle  a  delivery  tube  reaches  beneath  the 
surface  of  a  solution  of  sodium  hydroxide,  which  absorbs  any 


H20    H2S04 
FIG.  19. 


NaOH 


escaping  chlorine,  and  enables  us  to  know  when  the  air  in  the 
apparatus  has  been  displaced.  When  a  collecting  bottle  is  full 
of  chlorine  it  is  removed  and  stoppered,  and  replaced  by 
another  bottle  until  enough  gas  has  been  obtained.  The  bath 
of  hot  water  is  now  replaced  by  one  of  cold  water,  and  the 
evolution  of  chlorine  is  thus  stopped. 

The  proportions  by  weight  of  the  factors,  manganese 
dioxide  and  hydrochloric  acid,  and  of  the  products  re- 


OTUEE  METHODS.  81 

suiting  from  their  action  on  one  another  are  as  fol- 
lows :  — 

87  grams  manganese  dioxide  and  146  grams  hydrochloric 
acid  (in  aqueous  solution)  give  71  grams  chlorine,  126  grams 
manganous  chloride,  and  36  grams  water.  The  same  facts  are 
represented  by  the  equation, 

MnO2  +  4  HC1  =  MnCl2  -f  2  H2O  +  C12. 

It  is  probable  that  the  reaction  takes  place  in  two 
stages,  and  that  manganese  tetrachloride,  MnCl4,  is 
formed  first  and  then  breaks  down  into  manganous 
chloride  and  chlorine.  Equations  representing  these 

facts  are  :  — 

-—  •-—       -^^  » 

(1)  Mn02  -f  4  HC1  =  MnCl4  +  2  H2O. 
(2) 


The  final  result,  in  any  case,  is  that  shown  in  the 
equation  given  above. 

Instead  of  manganese  dioxide  and  hydrochloric  acid, 
a  mixture  of  common  salt,  manganese  dioxide,  and  sul- 
phuric acid  is  often  used.  The  result  is  approximately 
the  same,  however,  for  common  salt  and  sulphuric  acid 
give  by  their  action  hydrochloric  acid. 

82.  Other  Methods.  —  Manganese  dioxide  is  not  the 
only  substance  that  will  liberate  chlorine  from  hydro- 
chloric acid  ;  potassium  dichromate,  K9Cr9O7  ;  potassium 
chlorate,  KC1O8;  nitric  acid,  HNO3,  and  many  other 
substances  will  do  it. 


82  CHLORINE. 

But  potassium  chlorate  and  hydrochloric  acid  give,  besides 
chlorine,  an  explosive  oxide  of  chlorine.  Similarly  the  chlo- 
rine formed  from  the  action  of  nitric  acid  upon  hydrochloric 
acid  is  mixed  with  other  substances.  Hence  these  methods 
are  not  used  to  prepare  gaseous  chlorine. 

The  mixture  of  nitric  and  hydrochloric  acids  is  used 
extensively,  however,  under  the  name  aqua  regia  (== 
royal  water)  as  a  solvent  for  gold,  platinum,  and  other 
metals  not  readily  attacked  by  single  acids.  Aqua  regia 
• —  called,  also,  nitro-hydrochloric  acid  —  is  thus  only  a 
source  of  chlorine.  Metals  dissolved  in  it  are  converted 
into  chlorides. 

It  will  be  noticed  that  the  substances  which  react 
with  hydrochloric  acid  to  give  chlorine  are  all  oxidizing 
agents.  The  liberation  of  chlorine  in  all  the  methods 
described  is  thus  brought  about  in  practically  one  way, 
namely,  by  the  oxidation  of  the  hydrogen  of  hydrochlo- 
ric acid  to  water,  part  of  the  chlorine  being  set  free. 

There  is  a  process  for  the  manufacture  of  crude  chlorine  on 
a  large  scale  by  the  use  of  atmospheric  oxygen  as  the  oxidizing 
agent  ;  this  is  known  as  Deacon's  process.  By  this  method, 
hydrochloric  acid  gas  mixed  with  air  is  passed  over  heated 
bricks,  which  have  been  soaked  in  a  solution  of  copper  sulphate, 
or  of  copper  chloride,  and  then  dried.  In  some  way  —  we  do 
not  know  just  how  —  the  oxygen  is  able,  under  these  condi- 
tions, to  act  like  the  oxidizing  agents  mentioned  above. 

Another  method  for  the  production  of  chlorine  —  and 
one  that  may  in  time  displace  all  others  —  consists  in 


PHYSICAL   PROPERTIES.  83 

electrolyzing  a  concentrated  solution  of  hydrochloric  acid 
or  of  sodium  chloride.  The  chlorine,  which  is  electro- 
negative, appears  at  the  positive  electrode,  and  the 
metal  or  hydrogen  at  the  negative  electrode  (cf.  §  37). 

83.  Physical  Properties. — Chlorine  is  about  2% 
times  as  heavy  as  air,  and  35.5  times  as  heavy  as  hydro- 
gen. It  is  easily  soluble  in  water,  more  than  two  vol- 
umes of  the  gas  being  absorbed  by  one  of  water  at  the 
ordinary  temperature.  The  solution  —  called  chlorine 
water  —  possesses  many  of  the  properties  of  the  gas. 

In  a  warm,  saturated  solution  of  common  salt,  chlorine  is 
only  slightly  soluble  ;  it  may,  therefore,  be  collected  over  brine 
instead  of  under  air. 

If  chlorine  is  passed  into  iced  water  and  the  solution 
is  cooled  below  0®  C.,  a  crystalline  substance  separates 
out ;  this  is  a  compound  containing  144  parts  of  water 
to  71  of  chlorine,  and  therefore  represented  by  the 
formula  C12.  8  H2O.  It  is  called  chlorine  hydrate. 
Use  may  be  made  of  this  substance  to  condense  chlorine 
to  the  liquid  state. 

For  this  purpose,  the  crys- 
tals of  chlorine  hydrate  are 
dried  between  filter  papers  or 
on  unglazed  clay  plates,  and 
then  put  into  the  closed  limb 
of  a  tube  bent  as  shown  in 
Fig.  20.  The  open  end  is 
then  sealed.  The  end  of  the 
tube  containing  the  chlorine  hydrate  is  now  warmed  in  a 


84  CHLORINE. 

water  bath  to  30°  C.,  while  the  other  end  is  surrounded  by  a 
freezing  mixture  of  ice  and  salt.  After  a  short  time,  liquid 
chlorine  will  condense  at  the  drawn-out  end.  The  gas  has 
been  liquefied  by  its  own  pressure.  This  experiment  can  be 
carried  out  only  in  a  strong  tube  securely  sealed. 

Liquid  chlorine  is  an  article  of  commerce  ;  it  is  stored 
and  transported  in  iron  cylinders. 

84.  Chemical  Properties.  —  Chlorine  is  a  very  dan- 
gerous substance  to  inhale,  and  should,  therefore,  be 
generated  only  in  a  gas  chamber,  or  where  there  is  a 
good  draught.  If  it  has  been  taken  into  the  lungs, 
alcohol  or  ammonia  should  be  inhaled  to  counteract,  it. 

It  is  always  well  to  sprinkle  a  little  ammonia  water  about  in 
the  neighborhood  of  a  chlorine  generator. 

Chlorine  is  intensely  active  toward  many  other  ele- 
ments, forming,  by  direct  union  with  them,  the  chlorides. 
Many  substances  that  combine  with  oxygen  slowly,  or 
not  at  all  at  ordinary  temperatures,  unite  readily  with 
chlorine.  Powdered  antimony  and  copper  foil  (the  lat- 
ter must  be  hot)  glow  when  put  into  chlorine,  the  prod- 
ucts being  antimony  trichloride  (SbCl3)  and  cupric 
chloride  (CuCl2)  respectively.  Sodium,  tin,  magne- 
sium, and  phosphorus  all  give  corresponding  chlorides 
when  put  into  the  gas.  But  it  is  toward  hydrogen  that 
chlorine  shows  its  most  remarkable  behavior,  for  while 
the  two  gases  do  not  combine  at  all  in  the  dark,  and 
only  very  slowly  in  diffused  light,  yet  they  unite  with 
explosive  violence  in  sunlight. 


ACTION  OF  CHLORINE  AND  AMMONIA. 


85 


The  mixture  of  hydrogen  and  chlorine  may  also  be  exploded 
by  a  burning  match  or  by  the  electric  spark. 

Chlorine  shows  this  tendency  to  combine  with  hydro- 
gen not  only  when  the  hydrogen  is  in  the  free  state,  but 
also  when  it  is  united  with  other  elements.  As  illus- 
trations we  may  take  the  action  of  chlorine  toward 
water,  ammonia,  and  turpentine. 

85.  Action  of  Chlorine  and  Water. — The  aqueous 
solution  of  chlorine  may  be  preserved  for  a  long  time  if 
kept  cold  and  in  the  dark,  but  it  decomposes  rapidly  in  sun- 
light, giving  as  final  products  hydrochloric  acid  and  oxy- 
gen. The  equation  is  :  — 

2H2O  +  2  C12 »4HC1  + 

O2.     (See,  however, "  bleaching 
powder.")    If  the  decomposition 
of   chlorine  water   by  sunlight 
is    carried   out   in  a  long  tube 
(Fig.  21),  a  colorless   gas  will 
collect  in  the  upper  part  of  the 
tube.    The  gas  is  oxygen.   Much     , 
of  this  gas  will  also  be  found 
in  the  solution. 


Oxygen 


Chlorine  Water 


Chlorine  Water 


FIG.  21. 

86.  Action  of  Chlorine  and  Ammonia.  —  When  the 
hydrogen  of  ammonia  is  appropriated  by  chlorine,  hydro- 
chloric acid  and  nitrogen  are  formed,  as  represented  by  the 
equation, 

2  NH3  +  3  C12 »  6  HC1  +  Na. 


With  the  excess  of  ammonia  the   hydrochloric  acid  gives 


86  CHLORINE. 

ammonium  chloride,  NH4C1,  the  material  of  the  white  smoke 
seen  when  ammonia  and  chlorine  gases  come  together. 

87.  Action  of  Chlorine  and  Turpentine.  — The  behav- 
ior of  chlorine  and  turpentine  may  be  shown  by  immers- 
ing a  piece  of  filter  paper  in  warm  turpentine    and    then 
plunging    it  into   a  jar  of  chlorine.     The    turpentine  soon 
ignites,    and  burns  with  a  smoky  flame.     Turpentine    is  a 
compound  of    carbon  and  hydrogen,  and  the   chlorine,  by 
uniting  with  some  of  the  hydrogen,  but  not  with  the  car- 
bon, sets  the  carbon  free  in  the  form  of  a  dense,  black  smoke. 

88.  Uses  of   Chlorine.  —  Chlorine    is  used  in  large 
quantities  as  a  bleaching  and  disinfecting  agent,  and 
is  generally  made  for    these    purposes    from    bleaching 
powder,  or  "  chloride  of  lime."     Bleaching  powder  is  a 
white  substance  formed  by  the  action  of  chlorine  upon 
"slaked"  lime  (calcium  hydroxide),  and  is  easily  de- 
composed by  acids,  even  by  the  carbon  dioxide  of  the 
air,  with  evolution  of  chlorine.     Fabrics  to  be  bleached 
by  the  chlorine  process  are,  therefore,  immersed  in  a 
bath  of   dilute   acid,  and    then   in    one    of  chloride  of 
lime.     In  this  way  chlorine  is  set  free  in  immediate  con- 
tact with  the  coloring  matter  of  the  cloth,  and  bleaches  it. 

Chlorine  is  not  a  bleaching  agent  ordinarily,  unless  water  is 
present,  hence  it  is  likely  that  chlorine  itself  does  not  act  upon 
the  coloring  matter,  but  upon  the  water.  As  a  result,  oxygen 
is  probably  set  free  ;  and  it  is  oxygen,  and  not  chlorine,  that 
bleaches  the  cloth.  Since  ordinary  oxygen  is  not  able  to  effect 
this  change,  we  assume  that  oxygen  at  the  instant  of  its  libera- 
tion from  water  is  in  a  condition  different  from  that  in  which 


EXERCISES.  87 

we  ordinarily  find  it.  We  say  that  it  is  in  a  nascent  condition. 
The  reason  for  the  peculiar  behavior  of  an  element  in  its  nas- 
cent state  will  be  given  later. 

Compared  with  the  old  bleaching  process,  which  con- 
sisted in  exposing  the  fabric  to  the  oxidizing  agents  of 
the  air,  the  chlorine  method  is  of  course  very  rapid, 
but,  unfortunately,  the  bleaching  agent  used  too  often 
attacks  the  fiber  of  the  cloth  as  well  as  its  coloring 
matter.  Hence  delicate  materials,  such  as  the  better 
grades  of  straws,  laces,  silks,  and  woolens,  are  usually 
decolorized  by  sulphur  dioxide,  which,  although  it  does 
not  bleach  so  permanently  as  chlorine,  has  yet  the  ad- 
vantage of  acting  less  upon  the  fabric. 

The  action  of  chloride  of  lime  as  a  disinfectant  is 
similar  to  its  action  as  a  bleaching  agent :  nascent  oxy- 
gen is  formed,  and  this  destroys  the  micro-organisms  of 
the  surrounding  air. 

89.  Exercises. 

1.  How  many  grams  of  chlorine  can,  theoretically,  be  ob- 
tained by  the  electrolysis  of  50  grams  of  hydrochloric  acid  ? 

2.  How  many  grams  of  manganese  dioxide  are  required  to 
give  with  an  excess  of  hydrochloric  acid  10  grams  chlorine  ? 

3.  What  will  be  the  volume  of  40  grams  chlorine  Under  con- 
ditions at  which  1  liter  of  hydrogen  weighs  0.09  gram  ? 

4.  How   much   silver   chloride,    AgCl,   can   be   formed   by 
burning  54  grams  silver  in  chlorine  gas  ? 

5.  Calculate  the  per  cent  of  chlorine  in  sodium  chloride. 

6.  How  many  liters  of  chlorine  can  be  made  by  the  action 
of  25  grams  of  manganese  dioxide  with  an  excess  of  hydro- 
chloric acid  ?    (Assume  that  1  liter  of  chlorine  weighs  3  grams.) 


CHAPTER  VII. 
HYDROCHLORIC  ACID. 

90.  Existence. — Hydrochloric   acid   is   a   colorless, 
heavy  gas  which  fumes  in  moist  air  and  dissolves  readily 
in  water.     It  is  found  in  only  small  amounts  in  nature, 
e.  g.,  in  volcanic  gases  and  in  some  springs.     It  makes 
up  about  0.02  of  1^  of  the  gastric  juice. 

91.  Preparation:   Common  Laboratory  Method. — 

Hydrochloric  acid  may  be  prepared  readily  by  the  action 
of  sulphuric  acid,  H2SO4,  upon  common  salt,  NaCl. 
The  apparatus  is  shown  in  Fig.  22. 


FIG.  22. 

The  flask  contains  common  salt  and  sulphuric  acid  diluted 
with  half  its  volume  of  water,  and  cooled.      (Caution!  In  di- 


COMMERCIAL  MANUFACTURE  OF  THE  ACID.       89 

luting  sulphuric  acid  we  always  pour  the  acid  into  the  water.) 
A  hot-water  bath  serves  to  heat  the  flask.  The  first  bottle 
contains  concentrated  sulphuric  acid  to  dry  the  gas,  and  the 
second  bottle  is  the  collecting  vessel.  A  beaker  of  water 
collects  any  escaping  gas  and  shows  when  the  gas  in  the  col- 
lecting bottle  is  free  from  air. 

The  reaction  between  common  salt  and  sulphuric  acid 
takes  place  according  to  the  equation, 

NaCl  -f  H2SO4 >  ^TaIISO4  -f  IIC1 ; 

that  is  to  say,  58.5  grams  of  sodium  chloride  and  98 
grams  of  sulphuric  acid  give  120  grams  sodium  hydro- 
gen sulphate  and  36.5  grams  hydrochloric  acid.  If 
there  is  an  excess  of  sulphuric  acid,  the  sodium  chloride 
is  all  used  up,  and  the  white  solid  which  crystallizes  in 
the  generating  flask  when  the  latter  cools  is  sodium 
hydrogen  sulphate. 

92.  Commercial    Manufacture  of   the    Acid.  —  In 

the  first  stage  of  the  manufacture  of  sodium  carbonate 
(soda)  by  the  "  Le  Blanc  "  process,  sodium  chloride  is 
treated  with  sulphuric  acid,  with  the  results  illustrated 
by  the  equation  just  given.  Since,  however,  sodium 
sulphate  (Na2SO4)  and  not  sodium  hydrogen  sulphate  is 
the  product  wanted,  the  sodium  hydrogen  sulphate  is 
converted  into  sodium  sulphate  by  heating  it  with  more 
sodium  chloride  to  a  high  temperature.  The  reaction 
which  then  takes  place  is  represented  thus :  — 

INaCl  -f  NaHSO, »  Na2SO4  -f-  HC1. 

58.5  g.    120  g.  142  g.    36.5  g. 


90  HYDROCHLORIC  ACID. 

From  this  we  see  that  the  hydrogen  of  sulphuric  acid, 
like  that  of  water,  may  be  replaced  in  two  stages  ;  for 
while  at  low  temperatures  only  58.5  grams  of  sodium 
chloride  react  with  98  grams  of  sulphuric  acid,  at  a  high 
temperature  a  second  quantity  of  sodium  chloride,  equal 
to  the  first,  is  able  to  react,  and  thus  produces  a  second 
quantity  of  hydrochloric  acid.  This  will  be  more  evi- 
dent when  the  two  equations  are  written  together  :  — 

(1)  NaCl  -f  g  >  S04  -  »  ga  >  S04  +  HC1. 

(2)  NaCl  +     a  >  S04  -  »     »  >  S04  +  HC1. 


The  hydrochloric  acid  formed  as  a  by-product  in  the 
soda  manufacture  is  conducted  into  water,  and  the  solu- 
tion is  sold  as  commercial  hydrochloric  acid.  It  is  usu- 
ally somewhat  colored  by  slight  impurities. 

93.  Physical  Properties.  —  Hydrochloric  acid  is 
about  li  times  as  heavy  as  air,  and  18^  times  as  heavy 
as  hydrogen.  It  is  very  soluble,  505  c.c.  being  held  by 
1  c.c.  water  at  0°  C.  and  760  mm.  pressure.  The  con- 
centrated solution  of  the  pure  gas  in  distilled  water  at 
the  ordinary  temperature  is  the  "  chemically  pure  "  (c.  p.) 
reagent  hydrochloric  acid.  This  is  a  colorless  liquid 
of  specific  gravity  1.2  (Water  =  1).  The  concentrated 
solution  of  hydrochloric  a,cid  fumes  strongly  in  moist  air 
because  the  escaping  gas  condenses  some  of  the  water 
vapor  of  the  air. 


COMPOSITION   OF  HYDROCHLORIC  ACID. 


91 


The  dry  gas  can  be  converted  into  a  colorless  liquid  at  a  low 
temperature  and  great  pressure. 

94.  Volumetric  Composition  of  Hydrochloric  Acid. 

-  The  composition  of  hydrochloric  acid  may  be  demon- 
strated in  the  same  way  as  that  of  water,  viz.,  by  electrol- 
ysis.    If  an  electric    cur- 
rent of  sufficient  strength 
is  passed  through  a  concen- 
trated aqueous  solution  of 
the  acid,  hydrogen  is  pro- 
duced at    the  —  electrode, 
and  chlorine  at  the  -f-  elec- 
trode. 


Cl 


Zn 


FIG.  23. 


The  gases  collect  at  un- 
equal rates,  at  first,  because 
of  the  greater  solubility  of  the 
chlorine  ;  but  when  the  liquid 
has  become  saturated  with  both 

gases  the  hydrogen  and  the  chlorine  gather  in  the  collecting 
tubes  at  the  same  rate.     See  Fig.  23. 

Therefore  hydrochloric  acid  gas  must  be  composed  of 
hydrogen  and  chlorine  united  in  equal  proportions  by 
volume. 

The  same  fact  may  be  proved  synthetically,  for  if  a 
mixture  of  hydrogen  and  chlorine  be  exploded,  it  will 
be  found  that  equal  volumes  of  the  two  gases  have  dis- 
appeared. The  volume  of  hydrochloric  acid  formed 
will  be  equal  to  the  sum  of  the  uniting  gases.  These 
facts  may  be  represented  graphically  as  follows :  — 


92  HYDROCHLORIC  ACID. 

n       +       n      — >     nn 

1  vol.  hydrogen.          1  vol.  chlorine.      2  vols.  hydrochloric  acid. 

Note  that  this  case  is  different  from  that  of  water ;  for,  in 
the  production  of  2  volumes  of  steam,  2  volumes  of  hydrogen 
united  with  1  of  oxygen,  i.  e.,  3  volumes  of  the  mixed  gases 
gave  only  2  volumes  of  the  product. 

Since  the  weights  of  equal  volumes  of  hydrogen  and 
chlorine  are  about  as  1  :  35.5,  the  two  volumes  of  hydro- 
chloric acid  formed  by  their  union  should  be  about  36.5 
times  as  heavy  as  one  volume  of  hydrogen.  Hence  hy- 
drochloric acid  should  be  about  18.25  (=  36.5  ±  2) 
times  as  heavy  as  hydrogen.  This  is  actually  the  case. 
That  hydrochloric  acid  gas  gives,  when  decomposed, 
one-half  of  its  own  volume  of  hydrogen,  may  be  shown 
by  the  action  of  sodium.  For  con- 
venience the  sodium  is  diluted  by 
alloying  it  with  mercury. 

A  small  amount  of  the  resulting  sodium 
amalgam  is  put  into  a  long  measuring  tube 
(Fig.    24)   full  of  hydrochloric  acid   gas. 
The  open  end  of  the  tube  is  then  closed 
with  the  thumb,  and  the  tube  is  shaken 
/Sodium  Amalgam  vigorously.     When  the  thumb  is  removed 
thumb-  under  water,  some   of  the  water   rushes 

FIG.  24.  UP  m^°  the  tube  to  replace  the  chlorine. 

If  the  water  inside  and  outside  the  tube 
is  now  brought  to  the  same  level,  the  volume  of  the  residual 
hydrogen  will  be  found  to  be  approximately  one-half  that  oi 
the  hydrochloric  acid  taken. 


CHLORIDES.  93 

95.  Acid  Properties.  —  In  addition  to  the  properties 
already  given,  hydrochloric  acid  (i.  e.,  the  aqueous  solu- 
tion or  the  moist  gas)  has  the  characteristics  of  a  special 
class  of  substances  called  acids.  Acids  may  be  de- 
scribed roughly  as  having  a  sour  (acid)  taste,  the  ability 
to  change  certain  vegetable  colors,  e.  g.,  blue  or  purple 
litmus  to  red,  and  also  the  power  to  neutralize  the  prop- 
erties of  another  class  of  substances,  viz.,  the  bases. 
Thus  the  action  of  hydrochloric  acid  upon  sodium  hy- 
droxide (a  base)  gives  sodium  chloride  (common  salt) 
and  water.  Metals,  too,  react  with  acids,  forming  salts. 
Thus,  zinc  with  hydrochloric  acid  gives  zinc  chloride 
(a  salt)  and  hydrogen. 

These  reactions  are  shown  in  the  equations,  — 
NaOH  +  HC1  -  »  N"aCl  -f  H2O,  and 


40  36.5  58.5         18 

Zn  +  2  HC1  -  »  ZnCl2  +  H2. 
65         73  136          2 

The  properties  of  acids,  bases,  and  salts  will  be  considered  in 
the  next  chapter. 

Hydrochloric  acid  is  one  of  the  most  important  acids. 
It  is  made  on  a  large  scale,  and  is  used  in  enormous 
quantities. 

96.  Chlorides.  —  The  chlorides  may  all  be  considered 
hydrochloric  acid  with  its  hydrogen  replaced  by  a  metal  ; 
the  acid  itself  is  often  called  hydrogen  chloride.  The 
most  important  chlorides  have  been  given  in  §  80  ; 


94  HYDROCHLORIC   ACID. 

others  are  barium  chloride,  BaCl2,  silver  chloride,  AgCl, 
and  ferric  chloride,  FeCl3.  The  most  abundant  chlo- 
ride is,  of  course,  common  salt,  NaCl. 

The  chlorides  of  most  of  the  common  metals  are  soluble  in 
water.  Exceptions  are  silver  chloride,  AgCl,  and  mercurous 
chloride,  HgCl.  Lead  chloride,  PbCl2,  is  only  slightly  solu- 
ble in  cold  water,  but  more  readily  in  hot  water. 

When,  therefore,  solutions  of  salts  of  silver,  lead, 
and  mercury  (in  its  mercuroMS  condition)  are  treated 
with  a  solution  of  a  chloride,  the  chlorides  of  these 
metals  are  precipitated. 

97.  Exercises. 

1.  300  c.c.  of  hydrogen  and  250  c.c.  of  chlorine  were  mixed 
and  exploded.     What  was  the  product  ?     Its  volume  ?     Which 
of  the  gases  used  was  in  excess  ?     How  much  ? 

2.  Calculate   the   percentage    composition    of    hydrochloric 
acid? 

3.  How  many  grams  of  sodium  chloride  are  needed  to  yield, 
with  sulphuric  acid,  20  grams  hydrochloric  acid  gas  ? 

4.  How  many  grams  hydrochloric  acid  can  be  made  from  35 
grams  potassium  chloride,  KC1  ? 

5.  What  weight  of  sodium  chloride  is  necessary  to  produce, 
with  sulphuric  acid,  20  liters  of  hydrochloric  acid  gas  when  1 
liter  of  hydrochloric  acid  gas  weighs  1.63  grams  ? 


CHAPTER   VIII. 
ACIDS,    BASES,    AND   SALTS. 

98.  Acids.  —  In  all  of  our  study  of  Chemistry  we 
shall  have  to  deal  constantly  with  bodies  belonging  to 
the  classes  adds,  bases,  or  salts.  Let  us  first  consider 
some  of  'the  acids.  One  of  these,  viz.,  hydrochloric 
acid,  we  have  already  studied  at  some  length;  other 
important  acids  are  nitric  acid,  sulphuric  acid,  acetic  acid, 
arid  tartaric  acid.  Only  a  short  description  of  these 
will  be  given  here. 

Nitric  acid,  HNO3,  is  a  colorless  liquid.  It  ordinarily 
has  a  sharp  odor  and  is  very  corrosive  in  concentrated  form. 
It  turns  the  skin  yellow.  A  dilute  solution  of  nitric  acid  is 
sour,  turns  blue  litmus  and  neutral  litmus  pink,  decomposes 
carbonates,  e.  g.,  marble,  or  calcium  carbonate,  and  acts 
upon  many  metals. 

Sulphuric  acid,  H2SO4,  is  a  heavy,  oily  liquid  which  dis- 
solves in  water  with  the  evolution  of  much  heat.  It  chars 
organic  substances,  and  therefore  becomes  dark  colored 
when  exposed  for  a  time  to  the  dust  of  the  ah-.  Its  dilute 
solution  has  a  sour  taste,  and  acts  upon  litmus  and  carbon- 
ates as  nitric  and  hydrochloric  acids  do.  With  many  metals 
dilute  sulphuric  acid  gives  sulphates  and  hydrogen. 

Acetic  acid,  HC2H3O2,  is  a  colorless,  sharp- smelling 
liquid.  Like  the  other  acids,  it  has  a  sour  taste,  acts  upon 


96  ACIDS,    BASES,   AXD   SALTS. 

litmus,  and   decomposes    carbonates.     Vinegar  is    a  dilute 
solution  of  acetic  acid. 

Tartaric  acid,  H2C4H4O6,  is  a  white,  crystalline  solid, 
soluble  in  water.  Its  solution  has  properties  similar  to 
those  of  the  other  acids. 

All  of  the  acids  named,  except  hydrochloric  acid, 
may  be  looked  upon  as  water  joined  to  an  oxide.  The 
oxide  is  called  the  anhydride  of  the  acid. 

Thus  nitrogen  pentoxide,  N2O5,  is  the  anhydride  of  nitric 
acid,  for 

N2O5  -f  H2O  =  2  HNO3. 

Similarly,  sulphur  trioxide,  SO3,  is  the  anhydride  of  sulphuric 
acid. 

The  most  important  property  of  all  acids  is  their 
power  of  reacting  with  the  hydroxides  of  metals. 
When  a  solution  of  any  of  the  above  acids  is  treated 
with  a  solution  of  a  metal  hydroxide,  e.  g.,  sodium  hy- 
droxide, an  evolution  of  heat  takes  place,  and  if  the 
correct  amount  of  sodium  hydroxide  is  used,  the  taste 
of  the  acid,  its  power  to  change  litmus,  and  to  act  upon 
carbonates  and  metals,  will  all  disappear.  The  acid  has 
been  neutralized  by  the  sodium  hydroxide. 

99.  Bases.  —  The  general  properties  of  sodium  hy- 
droxide and  its  relation  to  water  have  already  been, 
given  (<?/.  §  46). 

Sodium  hydroxide,  NaOH,  is  a  white  solid  which  at- 
tracts moisture  and  carbon  dioxide  from  the  air,  and  thus 


BASES.  97 

becomes  converted  into  sodium  carbonate.  Its  solution 
changes  pink  and  neutral  litmus  to  blue,  feels  soapy  to  the 
touch,  and  has  a  bitter,  alkaline  taste. 

Potassium  hydroxide,  KOH,  resembles  sodium  hydrox- 
ide closely.  Its  aqueous  solution  has  properties  almost 
identical  with  those  of  sodium  hydroxide. 

Ammonium  hydroxide,  NH4OH,  is  not  known  in  the 
free  condition.  Its  aqueous  solution  smells  strongly  of  am- 
monia, owing  to  the  constant  evolution  of  this  gas,  but  its 
reaction  to  litmus,  etc.,  is  much  like  that  of  sodium  and 
potassium  hydroxides. 

Calcium  hydroxide,  Ca(OH)2,  is  a  white  solid  made  by 
adding  the  necessary  amount  of  water  to  quicklime,  CaO. 
Calcium  hydroxide  is  slightly  soluble  in  water ;  the  solution 
is  called  lime-water. 

^Barium  hydroxide,  Ba(OII)2,  is  much  more  soluble  than 
calcium  hydroxide.  Its  solution  is  called  baryta  water. 

All  of  these  hydroxides,  and  many  more,  are  grouped 
together  under  the  general  name,  bases.  The  most 
active  bases  are  sodium  and  potassium  hydroxides, 
which,  with  ammonium  hydroxide,  are  called  alkalies. 

As  has  already  been  stated  (c/.  §47),  hydroxides  may  be 
looked  upon  as  water  with  half  of  its  hydrogen  replaced  by  a 
metallic  element.  This  is  true  of  calcium  hydroxide,  having 

QTT 

the  formula  Ca^TT,  no  ^ess  than  of  sodium  hydroxide,  NaOH. 

OTT 
The  formula  Ca^yr,  however,  must  be  thought  of  as  derived 

f  TT         OTT  Oxi. 

from  2  H2O,  or  -3  TT  ~~    \TT.     Similarly,  the  formula  FeOH  or 
<  OH 


98  ACIDS,   BASES,   AND   SALTS. 

(HOH 
Fe(OH)3,  for  ferric  hydroxide,  is  derived  from  \  HOH  by  re- 

(HOH 
placing  half  of  the  hydrogen  there  represented  by  Fe. 

100.  Neutralization. — A  comparison  of  the  formu- 
las of  the  hydroxides  named  will  show  that  they  all  con- 
tain the  group  OH  (called  hydroxyl)  taken  one  or  more 
times.  On  the  other  hand,  the  formulas  of  the  acids 
show  that  the  acids  all  contain  hydrogen,  taken  once, 
twice,  etc.,  in  each  formula.  Furthermore,  the  neutrali- 
zation of  acids  by  bases  produces  salts  and  water,  as  is 
shown  by  the  following  equations :  — 

KaOH  +  HC1  =  NaCl  +  H2O. 

KOH  +  HC1  =  KC1  +  H2O. 

2  KOH  +  H2S04  =  K2S04  +  2  H2O. 

2  NH4OH  +  H2S04  =  (NH4)2S04  +  2  H2O. 

NH4OH  +  HN03  =  NH4K03  +  H20. 

Ca(OH)2  +  H2SO4  =  CaSO4  +  2  H2O. 

Ba(OH)2  -f  2  HC1  =  BaCl2  +  2  H2O. 

The  neutralization  of  properties  which  takes  place 
when  a  basic  hydroxide  and  an  afcid.  are  brought  to- 
gether thus  consists  in  the  union  of  the  hydrogen  of  the 
add  with  the  hydroxyl  of  the  base  to  form  water.  The 
metal  of  the  hydroxide  and  all  but  the  hydrogen  of  the 
acid  are  found,  on  evaporation  of  the  water,  combined 
in  the  resulting  salt. 

The  formula  of  an  acid  minus  the  replaceable  hydro- 
gen is  called  the  acid  radical. 


SALTS.  99 

i or.  The  Action  of  Oxides  with  Acids.  —  The  ox- 
ides of  the  metals,  like  the  hydroxides,  react  with  acids 
to  form  salts  and  water.  Thus,  the  equation 

CaO  +  2  HC1  =  CaCl2  +  H2O 

might  represent  the  reaction  between  calcium  oxide  and 
hydrochloric  acid,  just  as  the  equation 

Ca(OH)2  +  2  HC1  =  CaCl2  +  2  H2O 

represents  that  between  calcium  hydroxide  and  the  acid. 
The  only  difference  is  that  the  second  equation  shows 
twice  as  much  water  as  the  first.  .  This  is  because  the 
hydroxide  is  the  oxide  plus  water. 

It  is  probable,  however,  that  the  oxides  of  the  metals  react 
with  acids  only  in  the  presence  of  water.  If  this  is  so,  we 
must  assume  that  it  is  the  hydroxide  that  reacts,  and  not  the 
oxide.  The  action  of  calcium  oxide  upon  aqueous  hydrochloric 
acid  would,  therefore,  be  represented  thus:  — 

(1)  CaO  +  H2O  =  Ca(OH)2. 

(2)  Ca(OH)2  +  2  HC1  =  CaCl2  +  2  H2O. 

For  the  same  reason  the  action  of  sodium  oxide  upon  sul- 
phuric acid  would  be  represented  thus:  — 

(1)  Na2O  -f  H2O  =  2  ISTaOH. 

(2)  2  NaOH  -f  H2SO4  =  Na2SO4  -f  2  H2O. 

102.  Salts.  —  Salts  may  be  looked  upon  as  acids  in 
which  hydrogen  has  been  replaced  by  metals.  By  no 
means  all  of  the  hydrogen  of  many  acids  is  replaceable 
by  metals,  but  when  all  that  can  be  has  been  replaced,  the 
resulting  substance  is  called  a  normal  salt.  The  normal 


100  ACIDS,   BASES,   AND   SALTS. 

salts  formed  by  such  strong  acids  as  sulphuric  acid, 
nitric  acid,  and  hydrochloric  acid  with  such  electro- 
positive metals  as  sodium,  potassium,  and  the  group 
ammonium,  are  neutral  in  their  properties,  i.  <?.,  tLey  will 
not  turn  blue  litmus  red,  nor  red,  blue,  but  will  pro- 
duce in  sensitive  litmus  a  characteristic  lavender  color. 
The  taste  of  such  salts  resembles  that  of  common  salt. 

But  a  salt  may  be  normal  without  being  neutral.  Thus,  so- 
dium carbonate  is  a  normal  salt,  but  the  reaction  of  its  solution 
toward  litmus  is  that  of  a  base.  Ferric  sulphate  solution,  on 
the  contrary,  has  an  acid  reaction. 

103.  Acid  Salts.  —  In  many  cases,  the  hydrogen  of 
the  acid  is  replaceable  by  metals  in  two  or  more  stages 
(cf.  §  92)  ;  hence  there  may  be  two  or  more  salts  formed 
by  the  acid  with  the  same  basic  hydroxide.  The  following 
equations  illustrate  this  :  — 


NaOH  +  H2SO4  =  KaHSO4  +  H2O. 
2  NaOH  +H2S04  =  Na2SO4  +  2  H2O. 
KOH  +  H2C03  =  KHC08  +  H2O. 
2  KOH  +  H2C03  =  K2C03  +  2  H2O. 

The  substance  NaHSO4  is  sodium  hydrogen  sulphate, 
while  Na2SO4  is  sodium  sulphate.  Similarly,  KHCO3  is 
potassium  hydrogen  carbonate,  but  K2CO3  is  potassium 
carbonate. 

The  salts  in,  which  there  is  still  replaceable  hydrogen 
are  called  acid  salts.  Acid  salts  may  usually  be  con- 
verted into  normal  salts  by  the  addition  of  enough  of 
the  hydroxide  to  replace  all  the  replaceable  hydrogen, 


BASIC  SALTS.  101 

and,   conversely,   normal  salts   may  be  converted  into 
acid  salts  by  treatment  with  free  acid.     Thus  :  — 

NaHSO4  +  XaOH  =  Ka2SO4  +  H2O,  and 
KaHCO3  -j-  NaOH  =  Ka^CO3  +  H2O. 

So,  also,        Na2SO4  -f  H2SO4  =  2  NaHSO4,  and 
4-  H2CO3  =  2 


Solutions  of  acid  salts  have  usually  an  acid  reaction 
toward  litmus,  but  not  always  ;  for  the  reaction  may 
sometimes  be  alkaline,  as  in  the  case  of  sodium  hy- 
drogen carbonate  and  disodium  hydrogen  phosphate, 
Na2HP04. 

104.  Basic  Salts.  —  A  salt  may  be  considered  from 
two  points  of  view,  either  (1)  as  an  acid,  the  hydrogen 
of  which  has  been  replaced  by  a  metal,  or  (2)  as  a  hy- 
droxide with  its  hydroxyl  (OH)  replaced  by  a  non- 
metallic  element  or  by  an  acid  radical  (cf.  §  100). 
Thus  sodium  chloride,  NaCl,  may  be  looked  upon  as 
hydrochloric  acid,  HC1,  with  hydrogen  replaced  by  so- 
dium, or  as  sodium  hydroxide,  NaOH,  with  hydroxyl 
replaced  by  chlorine.  Similarly,  calcium  chloride, 
CaCl2,  is  calcium  hydroxide,  Ca(OH)2,  with  hydroxyl 
replaced  by  chlorine;  and  calcium  sulphate,  CaSO4,  is 
calcium  hydroxide  with  hydroxyl  replaced  by  the  acid 
radical  SO4.  Now,  just  as  in  sulphuric  acid,  H2SO4, 
the  hydrogen  may  be  replaced  half  at  a  time,  so  in  cal- 
cium hydroxide  the  hydroxyl  groups  may  be  substituted, 
theoretically,  in  two  stages.  We  might,  therefore,  have 


102  ACIDS,   BASES,    AND   SALTS. 

from  calcium  hydroxide  and  hydrochloric  acid  two  com- 

OTT  r^i 

pounds,    (1)   Ca^   ,  and  (2)  Ca^.     Just  as  we  call  a 

salt  which  still  contains  replaceable  hydrogen  an  acid 
salt,  so  we  call  one  having  replaceable  hydroxyl  groups 
a  basic  salt. 

To  illustrate  :  Just  as  phosphoric  acid,  H3PO4,  may 
have  two  acid  sodium  salts,  viz.,  NaH2PO4  and  Na2HPO4, 
so  bismuth  hydroxide,  Bi(OH)3,  might  have  two  basic 

nitrates,  viz.,   BiVr'2    and    Bi          ,  .     The  normal 


nitrate  is,  of  course,  Bi(NO3)3. 

It  is  to  be  noted,  however,  that  some  basic  salts  have,  appar- 
ently, a  more  complex  constitution. 

105.  Basicity  and  Acidity.  —  The  number  of  stages 
in  which  the  replaceable  hydrogen  of  an  acid  can  be 
substituted  by  metals  determines  the  basicity  of  the 
acid.  Thus,  hydrochloric  acid,  HC1,  nitric  acid,  HNO3, 
and  acetic  acid,  HC9H3O9,  are  monobasic  acids  ;  carbonic 
acid,  H2CO3,  and  sulphuric  acid,  H2SO4,  are  dibasic; 
while  phosphoric  acid,  HgPO4,  is  tribasic. 

In  the  same  way,  the  number  of  stages  in  which  the 
hydroxyl  groups  of  a  basic  hydroxide  might  be  replaced  by 
acid  radicals  or  by  non-metals  determines,  roughly,  the  acid- 
ity of  the  hydroxide.  Thus  sodium  hydroxide  and  potas- 
sium hydroxide  are  monacidic  bases  ;  calcium  hydroxide, 
Ca(OH)2,  and  barium  hydroxide,  Ba(OH).2,  are  diacidic  ; 
while  bismuth  hydroxide,  Bi(OH)3,  is  triacidic. 


NOMENCLATURE   OF  ACIDS.  103 

106.  Nomenclature  of  Acids.  —  Acids  consisting  of 
hydrogen  and  an  electro-negative  element,  e.  g.,  hydro- 
chloric acid,  HC1,  are  designated  by  the  names  of  both 
elements.  Thus  HBr  is  hydrobromic  acid,  and  H2S  is 
hydro  sulphur  ic  acid. 

Such  compounds  are  often  named  like  ordinary  compounds 
of  two  elements  (c/.  §  77)  ;  thus:  hydrogen  chloride,  hydrogen 
sulphide,  etc. 

The  salts  of  such  acids  are  called  chlorides,  bromides, 
sulphides,  etc.  Thus,  the  sodium  salt  of  hydriodic  acid 
is  called  sodium  iodide,  and  the  barium  salt  of  hydro- 
chloric acid  is  barium  chloride. 

Most  acids,  however,  consist  of  a  non-metallic  element 
united  with  hydrogen  and  oxygen,  e.  g.,  nitric  acid,  sul- 
phuric acid,  phosphoric  acid,  acetic  acid,  etc. ;  these  are 
given  the  name  of  the  non-metallic  element  with  the 
final  syllable  ic.  This  method  of  nomenclature  applies, 
however,  only  to  inorganic  acids;  organic  acids,  e.g., 
acetic,  tartaric,  citric,  picric,  etc.,  acids  are  quite  arbi- 
trarily named. 

When  there  are  two  acids  containing  the  same  three 
elements  in  different  proportions,  the  ending  of  the  acid 
containing  the  greater  proportion  of  the  non-metallic 
element  is  made  OUS,  while  the  ending  of  the  one  con- 
taining less  of  this  element  is  made  ic.  Thus,  nitrogen 
forms  with  hydrogen  and  oxygen  at  least  two  com- 
pounds which  are  acids ;  of  these  the  one  containing  the 
more  nitrogen  —  its  formula  is  HNO| — is  called  nitrous 


104  ACIDS,   BASES,   AND   SALTS. 

acid,  while  the  one  containing  the  smaller  proportion  of 
nitrogen  is  called  nitric  acid.  So  sulphur  forms  with 
hydrogen  and  oxygen  both  sulphurous  acid,  H2SO3,  and 
sulphuric  acid,  H2SO4. 

The  acids  formed  by  the  element  chlorine,  however,  give 
the  best  illustrations  of  nomenclature.  Chlorine  forms  with 
hydrogen  and  oxygen  four  acid  compounds,  the  compositions 
of  which  are  represented  by  the  formulas  : — 

HC10, 
HC1O2, 

HClOg, 

HC1O4. 

The  second  of  these  compounds,  HC1O2,  is  called  chlorous  acid  ; 
the  third,  chloric  acid  ;  the  first,  because  it  contains  less  oxygen 
than  chlorous  acid,  is  called  hypochlorous  acid  ("  hypo  "  signifies 
"under"  or  "lower  than")  ;  the  fourth,  because  it  has  more 
oxygen  than  chloric  acid,  is  called  perchloric  acid. 

Acids  containing  a  non-metallic  element  united  with 
hydrogen  and  oxygen  are  called  oxygen  acids. 

107.  Nomenclature  of  Salts.  —  The  salt  of  an  acid 
ending  in  ous  is  given  the  ending  ite,  the  prefix  Tiypo 
remaining  unaltered,  if  present. 

Thus,  the  sodium  salt  of  nitrous  acid  is  sodium  nitrite, 
NaNO2 ;  the  potassium  salt  of  chlorous  acid  is  potassium  chlorite, 
KC1O2 ;  and  the  calcium  salt  of  hypochlorous  acid  is  calcium 
hypochlorite,  Ca(OCl)2. 

The  salt  of  an  acid  ending  in  ic  is  given  the  suffix 
ate?  the  prefix  per  remaining  unchanged,  if  present. 


NOMENCLATURE   OF  BASES.  105 

Thus,  the  ammonium  salt  of  nitric  acid  is  ammonium  nitrate  ; 
the  barium  salt  of  chloric  acid  is  barium  chlorate,  Ba(ClO3)2  ; 
the  potassium  salt  of  permanganic  acid  is  potassium  permangan- 
ate, KMnO4. 

In  many  cases  there  are  two  salts  of  the  same  metal 
with  a  given  acid,  as,  for  example,  two  iron  sulphates, 
designated  by  the  formulas  FeSO4  and  Fe2(SO4)3, 
respectively.  To  distinguish  between  these  the  ending 
of  the  name  of  the  metal  is  changed  to  ous  in  the  case  of 
the  sulphate  having  the  greater  proportion  of  the  metal, 
and  to  ic  in  the  case  of  the  compound  having  the  smaller 
proportion  of  the  metal. 

Thus,  FeSO4  is  called  ferrous  sulphate,  just  as  FeCl2  is 
called  ferrous  chloride  (c/.  §  78)  ;  and  Fe2(SO4)3  is  called  ferric 
sulphate,  just  as  FeCl3  is  called  ferric  chloride. 

1 08.  Nomenclature  of  Bases.  —  The  name  of  a  basic 
hydroxide  contains  the  names  of  all  the  elements  of 
which  the  hydroxide  is  composed.  The  ending  is  ide, 
the  radical  OH  being  treated  as  an  element.  When 
there  are  two  basic  compounds  of  Jiydroxyl  with  the 
same  metal,  the  name  of  the  metal  in  the  hydroxide 
having  the  larger  proportion  of  the  metal  ends  in  ous, 
while  the  ending  of  the  metal  in  the  other  hydroxide 
is  ic. 

Thus,  we  have  cuprous  hydroxide,  Cu2(OH)2,  and  cupric 
hydroxide,  Cu(OH)2  ;  also  ferrous  hydroxide,  Fe(OH)2,  and 
ferric  hydroxide,  Fe(OH)3. 


106  ACIDS,   BASES,   AND   SALTS. 

109.  Exercises. 

1.  What  is  formed  when  a  solution  of  sodium  hydroxide  is 
neutralized   by  nitric  acid?     Barium  hydroxide  by  sulphuric 
acid  ?     Ammonium  hydroxide  by  acetic  acid  ? 

2.  Name  the  calcium  salt  of  carbonic  acid  ;  the  lead  salt  oi 
hydrochloric   acid ;   the   potassium    salt   of   chloric    acid ;   the 
barium  salt  of  hypochlorous  acid  ;  the  sodium  salt  of  chromic 
acid  ;  the  silver  salt  of  hyponitrous  acid. 

3.  8  grams  of  sodium  hydroxide  are  contained  in  50  c.c.  of 
a  solution  ;  how  many  grams  would  this  be  in  every  liter  ? 

4.  112  grams  of  potassium  hydroxide  are  required  to  neu- 
tralize all  the  hydrochloric  acid  in  a  solution  ;  how  much  of 
the  acid  was  there  ?     If   the  solution  were  evaporated,  what 
salt  would  be  found?     How  many  grams  of  it? 

5.  What  is  the  formula  of  the  acid  salt  formed  by  sodium 
hydroxide  and  sulphurous  acid  ?     Its  name  ? 

6.  49  grams  of  sulphuric  acid  were   required  to  redden  lit- 
mus in  a  solution  of  potassium  hydroxide  ;  how  much  hydroxide 
was  there  in  solution?     How  much  potassium  sulphate  was 
formed  ? 

7.  What  are  the  formulas  of  the  basic  chlorides  theoretically 
possible  from  a  consideration  of  the  formula  Bi(OH)3  ? 


CHAPTER   IX. 


NITROGEN   AND    THE   ATMOSPHERE. 

1 10.  Existence  of  Nitrogen. — Nitrogen  is  found 
uncombined  chiefly  in  air,  of  which  it  makes  up  about 
IS'/o  by  volume  and  75.5  c/o  by  weight.  It  is  found  com- 
bined with  many  elements,  as  with  hydrogen  in  ammonia, 
and  with  hydrogen  and  oxygen  in  nitric  acid.  Nitro- 
gen is  an  essential  constituent  of  all  animals  and  of 
many  plants. 


in.  Preparation.  —  Crude  nitrogen  may  be  prepared 
from  air  by  the  removal  of  the  oxygen  by  means  of 
phosphorus  or  copper.  With  phosphorus  the  operation 
is  as  follows :  — 

A  vessel  (Fig.  25)  of  air  is  placed  over  a  bit  of  burning 
phosphorus  which  is  floated  in 
a  small  dish  upon  water.  The 
phosphorus  unites  with  the  oxy- 
gen in  this  confined  portion  of 
air  to  form  phosphorus  pentoxide, 
which  is  a  white  solid  easily  dis- 
solved by  the  water.  If  the 
experiment  is  carried  out  accu- 
rately, the  water  which  rises  FIG.  25. 
into  the  vessel  after  the  experi- 
ment is  a  measure  of  the  oxygen  used  up  by  the  phosphorus. 

107 


I  !    \ 


108 


NITROGEN  AND    THE  ATMOSPHERE. 


The  oxygen  of  air  is  removed  more  satisfactorily  by 
hot  copper.  The  apparatus  is  shown  in  Fig.  26. 

Purer  nitrogen  may  be  prepared  by  heating  ammonium 
nitrite,  NH4NO2.  The  reaction  takes  place  as  shown 
by  the  equation,  — 


NH4N02  ==  N2 


2  H20. 


64 


28 


Instead  of  ammonium  nitrite,  a  mixture  of  solutions  of  am- 
monium chloride  (XH4C1)  and  sodium  nitrite  (NaNO2)  is  gen- 


Water 


FIG.  26. 


erally  used.     When  this  mixture  is  heated  carefully,  a  regular 
stream  of  nitrogen  is  evolved. 


112.  Properties  of  Nitrogen. — Pure  nitrogen  is  a 
gas  without  taste,  odor,  or  color,  and  about  0.97  as 
heavy  as  air.  100  c.c.  of  water  under  ordinary  condi- 
tions can  dissolve  only  about  1  c.c.  of  the  gas.  One 


CHARACTER  OF  THE  ATMOSPHERE.      109 

liter  of  nitrogen  weighs  about  1.25  grams  at  standard 
conditions,  i.  e.,  at  0°  C.  and  760  mm.  pressure. 

Since  ordinary  combustion  and  respiration  require 
oxygen,  it  naturally  follows  that  "  atmospheric  nitrogen," 
i.  e.,  air  deprived  of  oxygen,  no  longer  supports  either  com- 
bustion or  respiration.  Pure  nitrogen,  like  atmospheric 
nitrogen,  is  an  extremely  inactive  substance,  combining 
directly  with  only  a  few  elements.  It  does  combine 
with  magnesium,  titanium,  lithium,  etc.,  at  an  elevated 
temperature,  giving  nitrides. 

Under  the  influence  of  the  electric  spark,  nitrogen 
unites  with  hydrogen  and  oxygen  to  form  nitrous  and 
nitric  acids,  and  with -hydrogen  alone  to  form  ammonia; 
hence  these  compounds,  and  the  substances  formed  from 
them,  viz.,  ammonium  nitrite  and  ammonium  nitrate,  are 
found  in  the  atmosphere,  in  natural  waters,  and  in  the 
soil. 

113.  "  Atmospheric   Nitrogen  "  a  Mixture.  —  The 

nitrogen  obtained  by  removing  oxygen  from  air  was  long 
considered  pure,  but  careful  experiment  made  the  weight  of 
a  liter  of  atmospheric  nitrogen  1.2571  grams  and  that  of  a 
liter  of  pure  nitrogen  1.2^07  grams.  The  cause  of  this  dif- 
ference was  investigated  and  found  to  be  due  to  the  fact 
that  atmospheric  nitrogen  contains  another  substance  heavier 
than  nitrogen.  Thus  argon  was  discovered  in  1894. 

114.  Character  of  the  Atmosphere. --The   atmos- 
phere is  the  gaseous  mantle  of  the  earth.     Some  of  its 
ingredients  are  practically  constant  in  amount,  but  others 
are  variable. 


110  NITEOGEX  AND    THE  ATMOSPHERE. 

Constant  Ingredients.  Variable  Ingredients. 

Nitrogen,  Water, 
Oxygen,                                        v      Carbon  dioxide, 

Argon,  Ozone, 

Helium,  Hydrogen  peroxide, 

Hydrogen,  Ammonium  nitrite, 

and  several  rare  Dust, 

and  recently  discovered  etc. 
substances. 

Bacteria  are  so  universally  present  and  of  such  great  im- 
portance to  many  changes^taking  place  in  the  atmosphere  that 
they  may  rightly  be  classed  among  its  variable  ingredients. 

By  pure  air  we  mean  a  mixture  of  the  constant  con- 
stituents of  the  atmosphere.  The  proportions,  by  vol- 
ume and  by  weight,  of  the  three  most  abundant  of  these 
are  as  follows  :  — 

BY  VOLUME.  BY  WEIGHT. 

Nitrogen,  78.06%  75.5% 

Oxygen,    21.00%  23.2% 

Argon,         0.94%  1.3% 

Hydrogen  exists  in  small  quantities  in  the  atmos- 
phere. Recent  experiments  with  the  air  of  Paris  show 
that  100  liters  of  it  contain  about  19  c.c.  of  hydrogen. 
Hydrogen  is  thus  present  in  almost  as  great  an  amount 
by  volume  as  carbon  dioxide  (cf.  §  117). 

Nitrogen  and  oxygen,  the  most  abundant  constituents  of  air, 
have  been  described  already.  Argon  and  helium,  while  not  at 
all  comparable  with  the  former  two  elements  in  importance, 
are  interesting  because  they  have  only  recently  been  discovered 
in  the  earth  ;  hence  a  short  description  of  each  follows, 


HELIUM.  Ill 

115.  Argon.  —  The  discovery  of  argon  almost  took 
place  a  hundred  years  before  this  substance  was  actu- 
ally studied    by  Ramsay   and    Rayleigh    in    1894 ;  for 
Cavendish,  the  discoverer  of  hydrogen,  records  the  obser- 
vation that  he*  could  not  get  all  the  nitrogen  of  the  air 
to  combine  with  oxygen  by  "  sparking  "  a  mixture  of 
these    gases  in  the    presence   of  potassium  hydroxide. 
This  was  in  1785.    The  "  residual  nitrogen  "  was  argon 
and  the  other  inert  gases  which  are  mixed  with  atmos- 
pheric nitrogen.    By  repeating  Cavendish's  experiments, 
Ramsay  and  Rayleigh  obtained  argon. 

A  second  way  of  obtaining  this  substance  is  to  pass  pure 
air  over  heated  copper,  which  takes  up  the  oxygen,  and  then 
over  magnesium,  which  absorbs  the  nitrogen  as  magnesium 
nitride,  Mg3N2.  The  nitrogen  may  also  be  removed  by 
lithium  or  calcium. 

Argon  may  be  condensed  to  a  colorless  liquid,  boiling 
at  — 185°  C.,  and  at  lower  temperatures  may  even  be 
obtained  in  the  solid  state.  In  gaseous  form  it  is 
heavier  than  oxygen.  It  is  much  more  soluble  in  water 
than  nitrogen,  hence  in  air  which  has  been  dissolved 
in  water  and  afterward  expelled  from  solution,  the  pro- 
portion of  argon  is  greater  than  in  the  atmosphere. 
Argon  is  almost  without  chemical  activity,  hence  its  name. 

116.  Helium.  — By  means  of  the  spectroscope,  helium 
was  discovered  to  be  a  constituent  of  the  sun  a  quarter 
of  a  century  before  it  was  known  to  exist  on  the  earth. 


112  NITROGEN  AND    THE  ATMOSPHERE. 

It  has  been  found  in  small  amount  in  the  earth's  atmos- 
phere, in  certain  rare  minerals,  in  some  springs,  and  in 
a  meteorite,  as  well  as  in  the  atmospheres  of  the  sun  and 
certain  fixed  stars.  Like  argon,  helium  is  very  inert. 
It  is  probably  less  soluble  in  water  than  any  other  gas. 

117.  Carbon  Dioxide  in   the    Atmosphere. — The 

chief  variable  constituents  of  the  atmosphere  are  carbon 
.dioxide  and  steam,  and  to  the  presence  of  these  two 
substances  many  of  the  properties  of  the  atmosphere  are 
due.  The  atmosphere  supports  the  life  of  chlorophyll- 
producing  plants  largely  because  it  contains  carbon 
dioxide. 

The  presence  of  carbon  dioxide  in  the  atmosphere  may  be 
shown  by  drawing  a  current  of  ordinary  air  through  a  solution 
of  calcium  hydroxide  (lime-water)  ;  the  white  solid  which  sepa- 
rates from  solution  is  calcium  carbonate,  CaCO3.  Its  forma- 
tion indicates  the  presence  of  carbon  dioxide  in  the  air. 

Under  ordinary  conditions,  10,000  parts,  by  volume, 
of  air  contain  only  3  or  4  parts  of  carbon  dioxide.  If 
this  relative  amount  of  the  gas  is  doubled  as  a  result  of 
respiration,  the  air  becomes  foul,  not  so  much  because 
of  the  carbon  dioxide  itself  as  because  of  the  decaying 
organic  matter  which  is  exhaled  along  with  it  from  the 
lungs  of  animals.  The  quantity  of  carbon  dioxide  in 
the  atmosphere  of  a  room  thus  serves  as  an  index  to  the 
amount  of  poisonous  material  present. 

The  great  weight  of  the  earth's  atmosphere  may  be  illus- 
trated by  the  fact  that  its  carbon  dioxide,  although  so  small  a 


ATMOSPHERIC  DUST.  113 

proportion  of  the  whole,  is  estimated  to  weigh  over  five  thou- 
sand billions  of  tons. 

118.  Water  Vapor  in  the  Atmosphere.  —  The  quan- 
tity of  water  vapor  which  the  atmosphere  is  capable  of 
holding  at  any  given  time  depends  upon  the  tempera- 
ture and  the  pressure.     Air  is  saturated  with  water  vapor, 
or  at  the   u dew-point"  when  the   slightest  reduction  of 
temperature  or  increase  of  pressure  causes  precipitation 
of  some  of  the  water. 

One  hundred  liters  of  air  at  25°  C.  and  at  ordinary  pressures 
can  hold  a  little  over  2  grams  of  water.  If  the  temperature 
falls  to  0°  C.,  about  1.7  grams  of  the  water  are  precipitated  as 
rain.  Usually  the  atmosphere  is  far  from  having  all  the  water 
it  can  hold,  only  60%,  or  less,  of  this  amount  being  present  on 
a  fair  day.  If  there  is  much  more  than  this,  we  recognize  its 
presence  by  the  "  closeness  "  of  the  atmosphere. 

119.  Atmospheric    Dust.  —  The    importance    of    at- 
mospheric dust  in  causing  certain  phenomena  is  well 
known.     It  causes  sunset  arid  sunrise  colors,  and  helps 
to  effect  the  precipitation  of  water  vapor  as  clouds  and 
rain.     An  experiment  to  illustrate  the  latter  influence 
is  the  following :  — 

A  large  flask  (Fig.  27)  is  filled  with  dust-free  air  by  drawing 
through  it,  for  some  hours,  air  which  has  passed  through  along 
tube  of  cotton  wool.  If  the  room  is  now  darkened,  and  abeam 
of  light  is  directed  through  a  small  opening  toward  the  flask,  the 
beam  will  be  visible  in  the  outer  air,  but  not  in  the  flask,  owing 
to  the  absence  of  dust.  A  small  amount  of  steam  is  next  in- 
troduced into  the  flask  by  connecting  the  flask  at  a  with  a  ves- 


114  NITROGEN  AND    THE  ATMOSPHERE. 

sel  of  boiling  water,  opening  the  pinch  clamps  at  a  and  6,  and 
sucking  for  a  moment  with  the  aspirator.  A  beam  of  light, 
directed  as  before,  will  be  practically  invisible.  The  beam  will 
still  be  invisible  when  the  flask  is  connected  with  the  aspirator 
and  partly  exhausted.  A  small  funnel  is  now  attached  to  the 
tube  at  a,  and  a  quantity  of  smoke  is  produced  in  the  mouth 


Cotton  wool 

1 — i — ^(T^T^^ 

To  Aspiratoi 


FIG.  27. 

of  the  funnel.  This  smoke  is  sucked  into  the  large  flask  by 
opening  the  clamps  at  a  and  b,  and  thus  making  connection 
with  the  aspirator.  If  now  the  clamp  a  is  closed,  and  the  air 
is  slightly  exhausted  through  &,  a  beam  of  light  passing  through 
the  flask  will  be  clearly  visible  owing  to  the  small  drops  of 
water,  i.  e.,  mist,  which  fill  the  flask.  In  a  similar  way  the 
dust  of  the  air  probably  causes  the  formation  of  drops  of  water. 

120.  Weight  and  Pressure  of  the  Atmosphere.  - 
One  liter  of  air  weighs,  at   standard  temperature  and 


LIQUEFACTION   OF  AIR.  115 

pressure,  1.293  grams,  and  is,  therefore,  T}g-  as  heavy  as 
water,  and  about  14.4  times  as  heavy  as  hydrogen.  Be- 
cause of  its  weight,  the  atmosphere  exerts  pres- 
sure upon  all  bodies  immersed  in  it.  This 
pressure  is  measured  by  the  height  of  the 
column  of  mercury  the  atmosphere  will  sup- 
port in  the  barometer  (Fig.  28).  The  mean 
height  of  the  barometer  at  0°  C.  and  at  sea 
level  is  760  mm.;  this  is  the  standard  baro- 
metric height,  and  is  called  a  pressure  of  one 
atmosphere. 

FIG.  28. 
A  column  of  mercury  760  mm.  high  and  1  sq.  cm. 

in  cross  sectional  area  weighs  1,033.6  grams  ;  hence  this  is  the 
pressure  of  the  atmosphere  upon  1  sq.  cm.  at  sea  level.  If 
the  liquid  used  is  water,  which  is  -$y  as  heavy  as  mercury,  the 
height  of  the  barometer  will  be  13.6  times  as  great  as  with 
mercury,  i.  e.,  10.4  meters,  or  34  feet. 

I2i.  Liquefaction  of  Air.  —  The  condensation  of  air 
to  the  liquid  condition  is  brought  about  by  the  san:e 
methods  as  those  used  to  liquefy  other  "  permanent " 
gases,  and  differs  radically  from  the  liquefaction  of 
gases  like  ammonia,  chlorine,  and  sulphur  dioxide. 
These  gases  may  be  condensed  at  ordinary  temperatures, 
if  only  sufficient  pressure  is  applied ;  but  gases  like  air 
cannot  be  liquefied  at  the  ordinary  temperature  l>y  any 
pressure,  however  great.  Indeed,  pressures  up  to  3,600 
atmospheres  have  been  applied  without  avail.  This  is 

due  to  the  fact  that  there  is  for  every  gaseous  substance 

_ 


116  NITEOGEN  AXD   THE  ATMOSPHERE. 

a  maximum  temperature  above  which  the  gas  cannot  be 
liquefied;  this  is  called  the  critical  temperature  of  the  gas. 

A  "permanent"  gas  thus  differs  from  an  easily  condensible 
gas  in  this  respect,  viz.,  that  the  critical  temperature  of  a 
condensible  gas  is  above  the  ordinary  temperature,  while  that 
of  the  permanent  gas  lies  far  below  the  ordinary  temperature. 
Such  gases  as  air,  hydrogen,  etc.,  are,  therefore,  condensed 
only  at  a  very  low  temperature  and  great  pressure. 

Two  general  methods  are  used  to  liquefy  true  gases. 
In  the  first  method  the  gas  to  be  condensed  is  first  cooled 
to  its  critical  temperature,  and  is  then  subjected  to 
pressure.  In  the  second  method  the  gas  is  first  strongly 
compressed,  and  is  then  cooled  to  its  critical  tempera- 
ture. 

The  second  method  has  been  used  recently  to  liquefy 
air  on  a  commercial  scale. 

The  apparatus  used  consists,  essentially,  of  two  systems  of 
pipes,  the  pipes  of  the  outer  system  forming  a  jacket  surround- 
ing those  of  the  inner  system.  The  air  of  the  inner  system  is 
that  to  be  liquefied. 

By  means  of  compressing  engines  (c/.  Fig.  34),  air  is  forced 
into  each  system  under  great  pressure.  Much  heat  is  thereby 
evolved.  When  the  compressed  air  has  cooled  to  the  ordinary 
temperature,  some  of  the  air  of  the  outer  pipes  is  allowed  to 
escape.  The  sudden  expansion  of  the  air  remaining  in  the  outer 
pipes  then  causes  an  amount  of  heat  to  be  absorbed  which  is 
equal  to  that  evolved  when  the  compression  took  place.  The 
air  of  the  outer  pipes,  now  intensely  cold,  cools  the  compressed 
air  of  the  inner  pipes  below  the  critical  temperature  of  air,  and 
thus  liquefies  it. 


THE   PROPORTION   OF   OXYGEN  IN  AIR. 


117 


122.  Properties  of  Liquid  Air. — In  the  liquid  con- 
dition air  is  colorless,  has  about   the  same  density  as 
water,  and  boils  at  — 190°  C.,  under  ordinary  pressure. 
When  freshly  made,  liquid  air  is  about  half  oxygen,  but 
the  proportion  of  oxygen  increases  by  the  evaporation 
of    the   nitrogen   (liquid   nitrogen    boils   about    10°   C. 
below   liquid    oxygen)    until    the    liquid    is    over   90^? 
oxygen. 

Liquid  air  is  preserved  in  open,  double-walled  vessels  called 
Dewar  bulbs  (Fig.  29).     The    space    between    the    walls    of 
the  bulbs  is  exhausted  of  air  to  secure 
non-conductivity    of    heat.     Tin    or  CN     /O 

wooden  boxes  having  double  walls  filled 
with  silk  may  also  be  used. 

Alcohol,  liquid  carbon  dioxide, 
mercury,  etc.,  solidify  when  placed 
in  liquid  air,  and  steel  burns  in  it 
like  tinder ;  yet  the  hand  may  be 
held  in  it  for  a  short  time  with- 
out injury,  because  protected  by 
a  non-conducting  film  of  air  in  the 
gaseous  state. 

123.  Determination  of  the  Proportion  of   Oxygen 
in  Air.  —  The  amount  of  oxygen  in  a  given  volume  of 
air  is  usually  determined  either  (1)  by  absorbing  the 
oxygen,  or  (2)  by  exploding  the  air  with  a  known  vol- 
ume of  hydrogen. 

The   phosphorus  absorption  method,  a  crude    form  of 


118  NITROGEN  AND    THE  ATMOSPHERE. 

which  was  described  under  nitrogen  (<?f.  §  111),  is  car 
ried  out  more  accurately  as  follows  :  — 

A  gas  analysis  tube,  partly  full  of  air,  is  inverted  in  a  hydrom- 
eter jar  (Fig.  30),  and  the  gas  volume  is  carefully  measured. 
The    temperature    and   pressure  are  recorded.     A 
thin   stick  of  phosphorus   is   introduced   into   the 
tube  by  means  of  a  bent  wire,  and  is  left  there 
P.  twenty-four  hours.     The   phosphorus   is   then   re- 

moved, and  the  residual  gas  is  measured.  When 
the  necessary  corrections  have  been  made,  the 
volume  of  oxygen  absorbed  by  the  phosphorus  and, 
consequently,  the  per  cent  of  oxygen  in  the  origi- 
nal air  are  readily  calculated. 

The  absorption  of  the  oxygen  by  potassium 
pyrogallate  gives  similar  results. 

The  second   general 

method,  viz.,   the   ex-          .  ^O^^TT^N 
plosion  of  a  mixture  of  from  Spark  Coil 
FIG.  so.      known    volumes    of   air  "TTOTftT^ 

and  hydrogen,    may    be 
illustrated  as  follows:-  1Q  cc  ^ 

In  the  straight  eudiometer  tube 
shown  in  Fig.  31,  a  known  vol- 
ume of  air  is  mixed  with  an  ex- 
cess of  hydrogen,  and  the  electric 
spark  is  passed  through  the  mix- 
ture. All  of  the  oxygen  of  the  FIG  31 
air  taken  will  thus  unite  with  hy- 
drogen to  form  water.  The  volume  of  steam  formed 


THE  AIE  A   PHYSICAL  MIXTURE.  119 

will  be  equal  to  that  of  the  hydrogen  used  up,  but  the 
volume  of  liquid  water  produced  by  the  condensation 
of  the  steam  will  be  insignificant.  Hence  there  will  be 
a  contraction  of  volume.  One-third  of  this  contraction 
will  be  the  volume  of  oxygen  present  in  the  original  air. 

To  take  a  concrete  example  :  If  10  c.c.  of  each  of  the  gases, 
air  and  hydrogen,  are  taken,  the  volume  remaining  after  the 
explosion  will  be  about  13.7  c.c.,  i.e.,  the  contraction  will  be 
6.3  c.c.  Of  this  contraction  2.1  c.c.  (=J  of  6.3)  are  oxygen; 
hence  10  c.c.  of  air  contain  2.1  c.c.  (=21%). of  oxygen. 

The  proportions  of  the  chief  constant  constituents 
of  the  atmosphere  have  been  determined  in  many  differ- 
ent places,  and  have  been  found  everywhere  practically 
the  same.  This  is  a  remarkable  fact  when  we  remem- 
ber that  air  is  only  a  physical  mixture,  yet  it  is  a  neces- 
sary consequence  of  the  circulation  of  the  air  and  the 
rapid  diffusion  of  its  gases. 

124.  The  Air  a  Physical  Mixture.  —  That  air  is  not 
a  compound  of  nitrogen  and  oxygen  is  proved  by  several 
facts,  of  which  the  following  are  illustrations :  — 

(1)  When  nitrogen  and  oxygen  are  mixed,  there  is  no 
evidence  of  union,  such  as  evolution  or  absorption  of  heat, 
change  of  volume,  etc. 

(2)  If  air  were  a  chemical   compound,  like  the  nitrogen 
oxides,  it  ought  to  have  the  same  composition  in  the  liquid 
state  as  in  the  gaseous.     This,  however,  is  not  the  case,  as 
was  stated  in  the  description  of  liquid  air. 

(3)  There  is  no  reason  why  air,  if  a  compound,   should 


120  NITROGEN  AXD    THE  ATMOSPHERE. 

have  its  composition  changed  by  solution  in  water ;  yet  this 
is  the  case.  Air  that  has  been  expelled  from  solution  in 
water  contains  about  35%,  by  weight,  of  oxygen,  instead  of 
23%,  owing  to  the  fact  that  oxygen  is  much  more  soluble 
than  nitrogen. 

125.  Exercises. 

1.  Calculate  the  composition,  in  parts  per  cent,  of  a  mixture 
of  gases  76  c.c.  of  which  contain  48  c.c.  oxygen,  12.5  c.c.  hy- 
drogen, and  the  remainder  nitrogen. 

2.  It  has  been  calculated  that  in  certain  places  000  grams  of 
nitric  acid  fall  every  year  upon  an  acre  of  ground  :  can  you 
suggest  how  it  is  formed  ? 

3.  Why  is  not  all  the  carbon  dioxide  of  the  air  at  or  near  the 
earth's  surface  ? 

4.  What  is  the  height,  in  inches,  of  a  barometer  standing  at 
742  mm.?     (See  Appendix.) 

5.  16  c.c.  of  a  mixture  of  nitrogen  and  oxygen  were  put 
with    25  c.c.  pure  hydrogen,  and  exploded.     The   volume    of 
residual  gas  was  23  c.c.     How  many  cubic  centimeters  of  the 
original  mixture  were  oxygen  ? 

6.  How  many  grams  of  nitrogen  can  be  obtained  by  decom- 
posing 16  grams  of  ammonium  nitrite?     How  many  liters  at 
standard  conditions  ? 

7.  How  many  liters  of  "atmospheric  nitrogen  "  can  be  ob- 
tained from  10  liters  of  air  ?    What  will  be  its  weight  at  stand- 
ard conditions  ? 


CHAPTER   X. 
PROPERTIES  OF  GASES.     THE   MOLECULAR   THEORY. 

126.  Gases  and  Vapors.  —  The  word  "  gas  "  is  often 
used  to  denote  a  substance  which  is  in  the  gaseous  state 
under  ordinary  conditions,  while  "  vapor"  is  used  to 
mean  the  gaseous  condition  of  a  substance  which  is 
ordinarily  liquid  or  solid.  Thus,  we  speak  of  oxygen 
as  a  gas,  but  of  steam  as  the  vapor  of  water.  In  reality, 
however,  the  term  "gas"  can  properly  be  applied  only 
to  a  gaseous  body  above  its  critical  temperature.  In  this 
sense,  chlorine  and  hydrochloric  acid  are  not  gases 
under  ordinary  conditions,  while  nitrogen  and  hydrogen 
are.  Boyle's  and  Charles'  laws,  which  follow,  are 
applicable  in  a  strict  sense  only  to  true  gases.  They 
are,  however,  approximately  true  for  all  gaseous  bodies 
at  moderate  pressures. 


127.  Relation  of  the  Volume  of  a  Gas  to  Pressure. 

—  At  a  constant  temperature  the  volume  of  a  gas  is  in- 
versely proportional  to  the  pressure  the  gas  supports. 
This  is  Boyle's  law.  A  mass  of  gas  having  a  volume 
of  30  c.c.  at  700  mm.  thus  occupies  27.63  c.c.  at  760 
mm.  pressure,  and  35  c.c.  at  600  mm.  pressure,  as  is 
evident  from  a  solution  of  the  proportions,  — 

ttt 


122  PROPERTIES   OF  GASES. 

30  :  x  ::  760  :  700,  and 

30  :  x  ::  coo  :  700. 

The  product  of  the  number  representing  the  volume 
of  a  gas  and  the  number  representing  the  pressure  under 
which  the  volume  is  measured  is,  therefore,  a  constant ; 

or, 

v  X  p  =  constant. 

Since  increasing  the  pressure  will  diminish  the  volume 
of  a  gas,  it  will  plainly  increase  the  mass  of  gas  in  a 
given  volume,  i.  e.,  the  density.  Hence  Boyle's,  or,  as 
it  is  sometimes  called,  Mariotte's  law  may  be  stated 
thus :  The  density  of  a  gas  at  constant  temperature  is 
proportional  to  the  pressure. 

Problems  involving  changes  of  pressure  may  be  solved 
as  follows :  — 

Problem  1.  A  quantity  of  gas  measures  120  c.c.  when  the 
pressure,  as  indicated  by  the  barometer,  is  720mm.  ;  what  will 
its  volume  be  at  760  mm.  ? 

Solution.  Since  the  volume  varies  inversely  as  the  pressure, 
120  c.c.  in  changing  from  a  lower  to  a  greater  pressure  must 
diminish  in  volume.  Hence  the  proportion,  120  :  x  ::  760  :  720. 
Whence  x  =  the  volume  at  760  mm. 

Problem  2.  What  will  be  the  volume  at  650  mm.  of  a  quantity 
of  nitrogen  which  at  820  mm.  has  a  volume  of  50  c.c.  ? 

Solution.  Fifty  c.c.  in  changing  from  a  pressure  of  820  mm. 
to  one  of  650  mm.  must  increase  in  volume.  Hence  the  pro- 
portion, 50  :  x  ::  650  :  820. 

128.  Relation  of  Volume  to  Temperature.  —  The 

pressure  remaining  constant,  a  change  of  temperature 


EELAT1ON  OF    VOLUME   TO    TEMPERATURE.      123 

alters  the  volume  of  a  gas  so  that  273  c.c.  at  0°  C.  be- 
come 274  c.c.  at  1°  C.,  283  c.c.  at  10°  C.,  303  c.c.  at  30°  C., 
253  c.c.  at  — 20°  C.,  etc.  The  reverse  changes  are  also 
true,  i.  e.,  253  c.c.  at  —20°  C.  become  273  c.c.  at  0°  C.; 
and  303  c.c.  at  30°  C.  become  283  c.c.  at  10°  C.  and  273 
c.c.  at  0°  C.  These  changes  correspond  to  an  increase 
of  ^  (—  0.00366)  of  the  gas  volume  at  0°  C.  for 
every  degree  the  temperature  is  raised  above  0°  C.,  and 
a  like  diminution  for  every  degree  the  temperature  of 
the  gas  falls  below  0°  C.  At— 273°  C.,  therefore,  the 
volumes  of  all  gases  would  be  0.  This  temperature 
(: — 273°  C.)  is  called  the  absolute  zero.  All  centigrade 
readings  may  be  changed  to  absolute  readings  by  add- 
ing 273  to  them.  Thus,  +10°  C.  ==  283°  abs.  C.,  arid 
— 15°  C.  =  258°  abs.  C.  The  law  expressing  the  rela- 
tion of  volume  to  temperature  (Charles'  law)  may, 
therefore,  be  stated  thus :  Under  constant  pressure  the 
volume  of  a  gas  is  proportional  to  its  absolute  tempera- 
ture. 

It  follows  that  the  pressure  exerted  by  a  gas  whose 
volume  is  kept  constant  is  proportional  to  the  absolute 
temperature  of  the  gas. 

Problems  involving  changes  of  temperature  are  solved 
as  follows :  — 

Problem  1.  Suppose  that  the  temperature  of  a  gas  measuring 
80  c.c.  at  0°  C.  is  changed  to  20°  C.  ;  what  volume  will  the  gas 
have  at  20°  C.  ? 

Solution.  Since  273  c.c.  at  0°  C.  become  293  c.c.  at  20°  C., 
80  c.c.  will  increase  a  proportional  amount.  Therefore  we 


124  PROPERTIES   OF  GASES. 

write  the  proportion,  80  :  x  ::  273  :  293.  Whence  x  =  the 
required  volume. 

Problem  2.  If  60  c.c.  of  oxygen  at  30°  C.  have  the  tempera- 
ture lowered  to  — 10°  C.,  what  will  the  volume  become  ? 

Solution.  Since  303  c.c.  of  a  gas  at  30°  C.  become  263  c.c.  at 
— 10°  C. ,  60  c.c.  of  oxygen  will  diminish  proportionally.  There- 
fore, 60  :  x  ::  303  :  263. 

129.  Reduction  to  Standard  Temperature  and  Pres- 
sure. —  In  §  123  we  learned  that  the  oxygen  of  an  en- 
closed portion  of  air  will  be  completely  removed  by 
phosphorus.  From  the  volume  of  air  taken  and  the 
volume  of  residual  gas  obtained  the  proportion  of  oxygen 
may  be  calculated.  Accurate  results  cannot,  however, 
be  expected  from  this  experiment  unless  the  original 
and  final  volumes  are  comparable,  i.  e.,  are  at  the  same 
temperature  and  pressure.  It  is  possible  that  conditions 
will  be  the  same  after  twenty-four  hours ;  probably, 
however,  the  barometric  reading  and  the  temperature 
will  have  changed,  and  it  will  be  necessary  to  reduce 
the  volumes  read  to  the  same  temperature  and  pressure 
for  comparison.  The  volumes  of  yases  are  usually  com- 
pared at  standard  temperature  and  pressure,  i.  e.,  at 
0°  C.  and  760  mm.,  respectively. 

Problems  of  this  kind,  involving  changes  in  both 
temperature  and  pressure,  are  solved  by  stating  the  pro- 
portion for  each  case,  and  then  combining  the  two 
proportions. 

Problem  1.  What  will  be  the  volume,  under  standard  con- 
ditions, of  40  c.c.  of  oxygen  measured  at  700  mm.  and  25°  C.  ? 


REDUCTION   TO   STANDARD    CONDITIONS.      125 

Solution.   («)  For  the  temperature-change  alone,  — • 
Since  298  c.c.  (273  -f  25)  at  25°  C.  become  273  c.c.  at  0°  C., 
40  c.c.  will  diminish  proportionally.     Therefore, 

40  :  x  ::  298  :  273. 
(6)  For  the  change  in  pressure, — 

Since  the  pressure  is  to  increase  from  700  mm.  to  760  mm., 
the  volume  will  be  diminished.  Hence  the  proportion, 

40  :  x  ::  700  :  700. 
Combining  the  two  proportions,  we  have 

298  :  273  ; 
40  :*  ::   760:700;°r' 

298  X  760a;  =  40  X  273  X  700. 

Whence  x  =  the  volume  which  40  c.c.  measured  at  25°  C. 
and  700  mm.  would  occupy  at  0°  C.  and  760  mm. 

Problem  2.  In  the  experiment  referred  to  in  §  104,  viz., 
the  analysis  of  air  by  phosphorus,  a  student  obtained  the  fol- 
lowing figures  :  — 

Volume  of  air  taken  for  analysis  =  45.6  c.c. 

Temperature  =  19°  C.    ' 

Barometer  reading  (corrected)  =  730  mm. 

On  the  next  day  the  figures  were, — 

Volume  of  residual  gas  =  35.7  c.c. 
Temperature  =  21°  C. 
Barometer  reading  =  742  mm. 

Required  the  volume,  per  cent,  of  oxygen  absorbed. 

Solution.  We  first  reduce  the  volume  of  air  taken  to  stand- 
ard conditions  by  solving  for  x  in  the  proportion,— 

4  292  :  273.' 

0  '  x  ' '  760  :  730. 

This  done,  we  get  the  volume  of  the  residual  gas  at  standard 
conditions  by  solving  for  x  in  the  proportion,  — 


126  PROPERTIES   OF  GASES. 

7  294  :  273. 

'  '  x  ' '  760  :  742. 

In  this  way  we  obtain  comparable  volumes.  The  per  cent 
of  oxygen  in  the  air  taken  may  now  be  found  by  solving  for  x 
in  the  equation, — • 

Vol.  of  air  — •  Vol.  of  residual  gas 


Vol.  of  air 


=  x. 


130.  Correction  of  the  Barometric  Reading.  —  The 

standard  height  of  the  barometer  is  760  mm.  at  0°  C. ;  at 
other  temperatures  the  reading  must  be  corrected  for 
temperature.  Thus  at  20°  C.  the  mercury  column  is 
3  mm.  longer  than  at  0°  C. ;  hence  this  amount  must  be 
subtracted  from  the  barometric  reading  at  20°  C.  to  give 
the  reading  at  0°  C. 

Another  correction  is  often  made  when  gases  are  meas- 
ured over  liquids,  viz.,  a  correction  for  the  tension,  i.  e., 
pressure,  of  the  vapor  of  the  liquids  (cf.  §  44).  The 
pressure  of  this  vapor  opposes  the  atmospheric  pressure, 
hence  the  volume  of  a  gas  read  over  a  liquid  is  greater 
than  it  would  be  if  the  gas  were  free  from  the  vapor  of  the 
liquid,  i.  e.,  dry.  To  get  the  volume  which  a  moist  gas 
would  have  if  dry,  we  must  assume  the  gas  to  be,  not 
at  the  pressure  read  from  the  barometer,  but  at  the  baro- 
metric pressure  minus  the  pressure  of  the  vapor  of  the 
liquid. 

A  table  of  the  tension  of  the  vapor  of  water  at  several  tem- 
peratures appears  in  the  Appendix. 

The  correction  for  the  tension  of  aqueous  vapor  is  not  ap- 
plied when  two  gas  volumes  which  are  to  be  compared  (</. 


THE  MOLECULAR  THEORY.         127 

§  129,  Prob.  2)  are  both  measured  over  water  at  approximately 
the  same  temperature  (in  which  case  the  aqueous  tensions  are, 
evidently,  equal),  nor  on  a  very  "  close  "  day,  when  the  atmos- 
phere itself  is  nearly  at  its  dew-point. 

131.  The  Molecular  Theory.  —  The  physical  be- 
havior of  gases,  as  well  as  that  of  liquids  and  solids, 
and  the  laws  which  govern  chemical  action  are  best  ex- 
plained by  the  assumption  that  matter  consists  of  par- 
ticles, or  molecules.  Molecules  are  assumed  to  be 
separated  from  one  another  by  distances  which  are  very 
small,  but  which  are  yet,  especially  in  the  case  of  gases, 
very  great  as  compared  with  the  size  of  the  molecules 
themselves.  When  we  speak  of  distances  between  mole- 
cules we  mean  average  distances ;  for  molecules  are  not 
at  rest,  but,  on  the  contrary,  in  constant  motion.  The 
cause  of  the  motion  of  molecules  is  their  own  inherent 
energy.  The  direction  of  their  motion  is  along  straight 
lines  except  as  the  molecules  collide  with  one  another 
or  with  the  walls  of  the  containing  vessel. 

Because  of  frequent  collisions,  the  energy  of  one  par- 
ticle of  matter  is  probably  not  the  same  as  that  of  other 
particles  in  the  mass.  When  we  measure  the  degree  of 
heat,  i.e.,  the  temperature  of  the  mass,  we  get  only  the 
mean  temperature  of  the  particles  in  the  neighborhood 
of  the  thermometer.  There  may  be  many  particles  far 
above  the  mean  temperature,  as  well  as  many  far  be- 
low it. 

The  above  is  a  short  outline  of  the  "  Molecular ^ Theory." 
The  student  must  remember  that  the  existence  of  molecules  is 


128  PROPERTIES   OF  GASES. 

assumed,  and  that  the  simple  explanation  of  phenomena  which 
this  assumption  permits  cannot  have  the  same  force  as  the 
phenomena  themselves.  Yet  so  necessary  does  this  explana- 
tion seem  at  the  present  time  that  scientific  men  everywhere 
hold  it. 

132.  The  Physical  States  of  Matter.  — The  theory 
of  the  existence  of  molecules  explains  the  behavior  of 
matter  in  its  three  physical  states   of  gas,  liquid,  and 
solid. 

When  a  substance  is  in  the  gaseous  state,  the  energy 
of  its  particles  is  so  great  that  the  gas  tends  to  fill  any 
space  presented  to  it,  however  great.  Hence  a  gas  con- 
fined in  a  vessel  exerts  pressure  upon  the  Avails  of  the 
vessel. 

In  the  liquid  state  the  energy  of  the  particles  is  less 
than  in  the  gaseous  condition.  At  the  same  time,  there 
is  still  great  freedom  of  motion  of  the  particles  among 
themselves,  so  that  a  liquid  takes  the  form  of  the  vessel 
containing  it. 

When  a  substance  is  in  the  solid  condition,  its  mole- 
cules are  not  ordinarily  at  liberty  to  alter  their  relative 
positions.  As  a  result,  a  solid  has  a  definite  form. 

133.  Avogadro's  Hypothesis. — Another  assumption 
regarding  molecules  is  of  great  importance,  viz.,  that  in 
equal  volumes  of  all  gaseous  substances,  at  the  same  tem- 
perature and  pressure,  there  is  practically  the  same  number 
of  molecules. 

This  is  the  hypothesis  enunciated  by  Avogadro,  an  Italian 


EXPLANATION  OF  DIFFUSION.  129 

physicist,  in  1811,  and  bj  Ampere  in  1814.  It  was  originally 
assumed  to  explain  the  similarity  in  the  physical  properties  of 
all  gases,  and  now  seems  so  probable  that  it  serves  as  a  basis  of 
chemical  reasoning.  Only  by  some  such  assumption  can  the 
laws  of  Boyle  and  of  Charles  be  explained  in  a  reasonable  way. 

134.  Explanation  of  Diffusion.  —  Having  in  mind 
the  molecular,  or  particle,  theory  of  matter,  we  are  able 
to  explain  many  common  phenomena  in  an  intelligent 
manner.  At  the  present  time  we  shall  confine  our- 
selves to  diffusion,  or  solution  (cf.  §  55). 

In  the  case  of  gases,  diffusion  takes  place  to  an  unlim- 
ited extent,  because  the  molecules  are  so  far  apart  that 
there  is  practically  no  resistance  to  the  passage  of  other 
particles  between  them. 

An  illustration  is  the  mixing  of  the  gases  composing  the  air. 

In  the  case  of  liquids,  however,  there  is  considerable 
resistance  to  diffusion,  hence  the  amount  of  diffusion 
is  usually  limited.  Thus,  when  ether  and  water  are 
mixed,  ether  dissolves  in  water,  and  water  in  ether,  up 
to  a  certain  limit. 

Solids  diffuse  into  'other  solids  very  slowly,  and  then 
only  ujider  great  pressure.  An  interesting  case  of  solid 
diffusion  is  the  formation  of  the  alloy  solder  from  its 
constituents  without  melting. 

In  Spring's  experiments,  lead  and  tin  were  subjected  to  a 
pressure  of  about  six  thousand  atmospheres  in  steel  molds. 
Under  these  conditions  the  solids  diffused  into  each  other,  as 
gases  do  at  ordinary  pressure.  The  solder  thus  made  has  the 
same  properties  as  that  obtained  by  fusion. 


130  PROPERTIES   OF  GASES. 

135.  Diffusion  of  Gases  and  of  Liquids.  —  When 
layers  of  gases  are  brought  over  each  other,  the  gases 
do  not  permanently  adjust  themselves  according  to  their 
relative  densities,  but  mix  with  each  other  until  any  given 
volume  of  the  mixture  contains  the  same  amount  of 
each  gas  as  every  other  equal  volume.  Thus,  if  a  jar 
of  hydrogen  is  placed,  mouth  downward,  over  a  jar  of 
carbon  dioxide  turned  mouth  upward,  and  the  junction 
between  the  jars  is  made  dose,  some  of  the  carbon  di- 
oxide will  ascend,  against  gravity,  into  the  upper  jar,  and 
some  of  the  hydrogen  will  descend  into  the  lower  one. 

This  is  due  to  the  fact  that  the  particles  of  carbon  dioxide 
having  the  greatest  energy  detach  themselves  from  the  layer 
of  this  gas  and  enter  that  of  the  hydrogen.  In  the  same  way, 
those  particles  of  hydrogen  which  have  the  greatest  energy 
leave  their  proper  layer  and  enter  that  of  the  carbon  dioxide. 
Thus  a  mixture  is  formed,  first  at  the  plane  of  contact  between 
the  two  layers,  but  gradually  extending  through  the  whole 
mass  of  the  gases. 

There  is  no  reason  to  suppose  that  diffusion  ever  ceases  in 
such  a  mixture. 

The  rate  at  which  diffusion  takes  place  is  very  differ- 
ent in  the  case  of  different  pairs  of  gases.  This  may  be 
illustrated  by  the  diffusion  of  bromine  vapor  in  air  and 
in  hydrogen.  If  a  drop  of  bromine  is  placed  in  each  of 
two  small  watch  glasses,  and  one  watch  glass  is  covered 
with  a  tall  jar  of  hydrogen  and  the  other  with  a  jar  of 
air,  bromine  will  rise  in  both,  against  gravity,  but  much 
more  rapidly  in  one  than  in  the  other. 


SOLUTION  OF  SOLIDS  IN  LIQUIDS.  131 

Diffusion  takes  place  in  the  same  way  between  two  liquids, 
and  between  solvents  and  solutions,  but  it  is  slower  in  these 
cases  than  with  gases. 

136.  Solution  of  Gases  in  Liquids.  —  When  a  layer 
of  a  gas  is  brought  over  a  liquid,  e.g.,  air  over  water,  gas 
particles  dissolve  in  the  liquid,  and,  on  the  other  hand, 
gas  particles  that  were  dissolved  leave  the  liquid,  until 
a  condition  of  equilibrium  is  reached  at  which  as  many 
gas  particles  enter  the  liquid  in  a  given  time  as  leave  it 
in  the  same  time. 

Similarly,  the  energy  of  some  of  the  particles  of  the 
liquid  causes  them  to  leave  the  liquid  surface,  and  to 
enter  the  gas  above,  until  as  many  particles  of  the  vapor 
of  the  liquid  return  to  the  liquid  in  a  given  time  as  leave 
it  in  the  same  time. 

If  the  gas  layer  above  the  liquid  is  in  motion,  as  when 
a  current  of  air  passes  over  water,  this  condition  of 
equilibrium  is  never  reached;  hence  the  water  evapo- 
rates. 

The  explanation  of  efflorescence  and  deliquescence  (cf.  §  53) 
is  found  here.  The  efflorescence  of  a  solid  is  due  to  the  fact 
that  the  crystal-water  of  the  solid  has  a  greater  vapor  tension 
((/.  §§  44  and  130)  than  that  of  the  water  of  the  atmosphere. 
Deliquescence,  on  the  contrary,  is  due  to  the  fact  that  the  water 
which  the  deliquescent  substance  has  absorbed  from  the  air 
has  a  smaller  vapor  pressure  than  that  of  the  water  of  the  air 
Hence  moisture  is  added  to  the  substance. 

137.  Solution  of  Solids  in  Liquids.  —  A   solid  dis- 
solves in  a  liquid  as  a  gas  does.     The  energy  of  the 


132  PROPERTIES   OF  GASES. 

particles  of  the  solid  causes  them  to  leave  the  surface  of 
the  solid  and  to  enter  the  solvent.  At  the  same  time 
there  is  a  constant  return  of  particles  that  have  dis- 
solved, to  the  solid  state.  These  two  operations  go  on 
until  there  are  as  many  dissolved  particles  becoming 
solid  in  a  given  time  as  there  are  solid  particles  being 
dissolved  in  the  same  time. 

When  this  condition  of  equilibrium  has  been  reached, 
the  solution  is  saturated,  not  because  no  further  diffu- 
sion between  solid  and  liquid  goes  on,  but  because  fur- 
ther action  cannot  affect  either  the  mass  of  undissolved 
solid  or  the  concentration  of  the  solution. 

The  gradual  solution  of  a  solid  may  be  illustrated  by  putting 
some  solid  potassium  permanganate,  KMnO4,  into  a  deep  bot- 
tle of  water  and  leaving  the  vessel  covered  and  undisturbed  for 
some  time. 

Since,  in  the  solution  of  a  solid  in  a  liquid,  the  par- 
ticles that  get  away  from  the  solid  are  those  having  the 
greatest  energy,  the  addition  of  heat  should  aid  solution 
by  giving  more  particles  sufficient  energy  to  enable 
them  to  leave  the  solid.  This  is  the  case  ;  for  the  solu- 
bility of  most  solids  increases  with  rise  of  temperature. 
If  the  heat  required  is  not  furnished  from  an  outside 
source,  it  comes  from  the  solvent.  Hence  the  dissolv- 
ing of  solids  usually  causes  a  reduction  of  temperature. 

138.  Osmotic  Pressure.  —  We  can  observe  and 
measure  the  tendency  of  the  particles  of  a  dissolved 
substance  to  diffuse  through  a  solution  by  preventing 


LAWS  OF   OSMOTIC  PRESSURE.  133 

the  diffusion.  To  do  this,  we  use  a  cell  which  permits 
the  solvent  to  pass  through  its  pores,  but  does  not 
allow  the  dissolved  substance  to  get  through.  Such  a 
cell  is  said  to  be  semi-permeable.  The  pellicles  of 
most  organic  cells,  many  animal  membranes,  and  certain 
inorganic  substances,  e.  g.,  copper  ferrocyanide,  have  this 
property  of  semi-permeability. 

To  illustrate :  If  a  solution  of  sugar  in  water  fills  a  semi- 
permeable  cell,  and  the  cell  is  attached  by  a  water-tight 
joint  to  a  long  glass  tube  and  then  immersed  in  pure  water, 
there  will  be  a  very  noticeable  rise  of  the  solution  in  the 
glass  tube.  The  height  to  which  the  liquid  in  the  tube 
rises  above  the  level  of  the  water  outside  the  cell  indicates 
how  much  the  pressure  inside  the  cell  exceeds  the  atmos- 
pheric pressure  at  the  time  of  the  experiment.  This  excess 
of  pressure  is  due  to  the  sugar  particles  in  the  cell,  and  is 
called  the  osmotic  pressure  of  the  solution. 

What  happens  when  the  cell  of  sugar  solution  is  im- 
mersed in  pure  water  is  that  water  from  without  enters 
the  cell  until  the  pressure  of  the  water  within  and  with- 
out is  the  same,  viz.,  one  atmosphere.  The  sugar  parti- 
cles, however,  not  being  able  to  diffuse  out  through  the 
semi-permeable  wall  of  the  cell,  add  their  pressure  to 
that  of  the  water,  and  force  the  liquid  up  the  tube. 

139.  Laws  of  Osmotic  Pressure.  —  If  the  strength 
of  the  solution,  i.  e.,  the  number  of  particles  of  dissolved 
substance  in  a  given  volume  of  the  solution,  is  increased, 
the  osmotic  pressure  increases  proportionally.  Since  the 


134  PROPERTIES  OF  GASES. 

strength  of  a  solution  means  the  same  as  density  in  the 
case  of  gases,  we  see  that  the  osmotic  pressure  of  a  solu- 
tion corresponds  to  gaseous  pressure  ;  for  the  latter,  as  we 
learned  in  §127,  is  proportional  to  the  density. 

Moreover,  the  osmotic  pressure  of  a  solution  increases 
proportionally  to  the  absolute  temperature,  as  in  the  case 
of  gaseous  pressure  (<?/.  §  128). 

Substances  in  dilute  solution  have,  therefore,  many  of  the 
properties  which  they  would  have  if  they  were  gases  and  had 
a  volume  equal  to  that  of  the  solution. 

140.  Exercises. 

1.  What  will  be  the  volume  at  0°  C.  of  a  mass  of  gas  which 
at  20°  C.  has  a  volume  of  102  c.c.? 

2.  How  many  c.c.  of  a  gas  measured  at  28°  C.  are  required 
to  fill  a  50  c.c.  eudiometer  at  0°  C.? 

3.  One   liter   of   air  at  15°  C.  will  have   what  volume   at 
25°  C.? 

4.  A  mass   of   air  measures  70  c.c.   when  the   barometric 
height  is  740  mm.     If  the  temperature  remains  constant,  what 
volume  ought  the  air  to  have  when  the  barometer  reading  is 
760  mm.?     When  it  becomes  650  mm.? 

5.  What  will  be  the  volume  at  800  mm.  of  a  quantity  of 
nitrogen  which  at  730  mm.  measures  1.2  liters? 

6.  A  quantity  of  hydrogen  which  measures  45  c.c.  at  18°  C. 
and  730  mm.  will  have  what  volume  at  0°  C.  and  760  mm.? 

7.  90  c.c.  of  air  at  20°  C.  become,  when  heated,  150  c.c. 
To  what  temperature  was  the  air  heated? 

8.  A  gas  measures  20  liters  at  10°  C.  and  760  mm.  pressure, 
hat  volume  will  it  have  at  30°  C.  and  720  mm.? 

,  The  weight  of  a  liter  of  nitrogen  at  standard  conditions 
about  1.25  grams,  that  of  1  liter  of  hydrogen,  .09  grains^ 


EXEECISES.  135 

and  that  of  1  liter  of  chlorine,  3.18  grams,  what,  upon  the  basis 
of  Avogadro's  hypothesis,  are  the  relative  weights  of  the  mole- 
cules of  nitrogen,  hydrogen,  and  chlorine  ? 

10.  If  in  Problem  9  we  call  the  weight  of  the  hydrogen 
molecule  2,  what  will  be  the  relative  weights  of  the  molecules 
of  nitrogen  and  chlorine  ? 


CHAPTER   XL 
AMMONIA. 

141.  Existence*  —  Ammonia  is  found  in    nature  in 
very  small  quantities.     As   already  stated  (cf.  §  112) 
electric  discharges  in  moist  air  produce  ammonia,  nitrous 
acid,   and   nitric   acid ;    consequently   ammonia   and  its 
compounds,  ammonium  nitrite  and  ammonium  nitrate, 
are  found  in  the  atmosphere  and  in  rain  water. 

Small  amounts  of  ammonium  compounds  are  found 
in  plants  and  in  animal  secretions. 

142.  Laboratory  Method  of  Preparation.  —  In  the 
laboratory,   ammonia  gas  may  be  made  either  (1)  by 
heating  an  ammonium  salt  with  slaked  lime,  or  (2)  by 
warming  concentrated  ammonium  hydroxide  solution. 

The  first  method  is  generally  carried  out  as  follows  :  — 

About  equal  parts  of  an  ammonium  salt,  e.  </.,  the  chloride  or 
the  sulphate,  and  slaked  lime  are  mixed  together  in  a  flask,  just 
enough  water  being  added  to  make  a  thick  paste.  The  flask 
is  connected,  as  shown  in  Fig.  32,  with  bottles  containing  water. 
None  of  the  connecting  tubes,  except  the  last,  dips  below  the 
surface  of  the  water. 

When  the  generating  flask  is  heated,  the  evolved  ammonia 
charges  the  water  of  the  collecting  bottles,  forming  ammonia 
water.  The  first  bottle,  especially,  will  show  a  rise  of  tempera- 
ture and  an  increase  of  volume. 

136 


LABORATORY  METHOD  OF  PREPARATION.    137 

If  the  gas  evolved  from  the  flask  is  passed  through 
bottles  or  U-tubes  filled  with  quicklime  or  with  solid 
sodium  "hydroxide,  it  may  be  used  for  studying  the  prop- 
erties of  dry  ammonia. 


FIG.  32. 

The  second  method  for*  producing  ammonia  gas  is 
more  convenient.  It  consists  simply  in  heating  a  small 
amount  of  concentrated  ammonium  hydroxide  solution. 
Much  ammonia  gas  is  evolved,  and  may  be  dried-  in  the 
usual  way. 

Calcium  chloride  ma}^  not  be  used  to  dry  ammonia,  for  the 
reason  that  ammonia  and  calcium  chloride  form  a  crystalline 
compound  of  the  formula  CaCl2.  8 


138  AMMONIA. 

The  reaction  between  slaked  lime  and  ammonium 
chloride  is  represented  by  the  equation,  - 

2  NH4C1  +  Ca(OH)2 »  CaCl2  +  2  NH4OH. 

The  ammonium  hydroxide  breaks  up  readily  into  am- 
monia and  water,  — 

NH4OH »  NH3  +  H20. 

143.  Commercial  Sources  of  Ammonia.  —  The  two 
chief  sources  of  ammonia  are  :  — 

(1)  The    dry    distillation    of  animal   matter    (bones, 
skin-products,  etc.). 

(2)  The  dry  distillation  of  soft  coal. 

Formerly  all  ammonia  and  all  ammonium  compounds 
were  made  by  the  dry  distillation  of  animal  matter; 
hence  the  name  " spirits  of  hartshorn,"  "spirit"  mean- 
ing a  distillate,  and  "  hart "  being  a  general  term  for 
"  deer,"  etc.  The  horns,  bones,  hoofs,  hides,  hair,  etc., 
of  animals  contain  nitrogenous  bodies,  which  give  much 
ammonia  when  heated  with  quicklime  without  access  of 
air. 

This  method  is  used  in  many  places,  especially  since  the  old 
method  of  making  illuminating  gas  has  been  abandoned,  to 
convert  into  a  valuable  product  almost  worthless  scraps  of  ani- 
mal matter. 

The  most  common  source  of  ammonia  for  many  years 
has  been  the  wash  liquors  of  the  gas  works.  Illumina- 
ting gas  is  made  in  the  old  process  by  the  dry  distillation 
of  soft  coal.  At  first  it  contains  many  impurities,  among 


PHYSICAL   PROPERTIES.  139 

them  ammonia,  hydrogen  sulphide,  and  carbon  dioxide  ; 
these  are  removed  by  passing  the  gas  through  water. 
When  the  wash  water,  containing  ammonium  sulphide, 
ammonium  carbonate,  etc.,  is  heated  with  slaked  lime, 
ammonia  gas  is  evolved. 

Ammonia  comes  into  the  market  as  ammonia  water 
(ammonium  hydroxide),  as  ammonium  salts,  and  also  in 
the  liquid  state.  To  produce  ammonium  salts  we  treat 
ammonia  water  with  acids  until  the  solution  is  neutral, 
and  then  evaporate  the  water. 

Thus,  ammonium  chloride  (sal  ammoniac)  results  when  the 
neutralized  solution  of  ammonium  hydroxide  and  hydrochloric 
acid  is  evaporated  ;  ammonium  nitrate  and  sulphate  are  formed 
if  nitric  acid  and  sulphuric  acid,  respectively,  are  used. 

The  equations  representing  the  formation  of  these  salts  are 
as  follows  :  — 


H  -k  HC1  =  NH4C1  +  H20. 
NH4OII  -j-  HNO3  =  NH4NO3  +  H2O. 
2  NH4OH  +  H2S04  =  (NH4),S04  +  2  H2O. 

144.  Physical  Properties.  —  Ammonia  is  a  colorless 
gas  of  characteristic  odor.  One  liter  of  it  under  stand- 
ard conditions  weighs  0.762  grams  ;  ammonia  is  there- 
fore 0.59  as  heavy  as  air. 

Ammonia  gas  is  very  soluble  in  water,  1  c.c.  of  water 
absorbing,  at  standard  conditions,  1,146  c.c.  or  0.873 
gram  of  the  gas.  The  specific  gravity  of  a  35J&  solu- 
tion at  15°  C.  is  about  0.88;  that  of  the  so-called  28° 
Beaume  solution  is  about  0.90. 


140  AMMONIA. 

Dry  ammonia  rushes  into  cold  water  as  into  a  vacuum,  and 
is  likewise  very  soluble  in  alcohol  and  in  ether.  Charcoal  can 
take  up  under  favorable  conditions  about  ninety  times  its  own 
volume  of  ammonia  gas. 

145.  Liquefaction  of  Ammonia.  —  The  critical  tem- 
perature of  ammonia  is  far  above  the  ordinary  tempera- 


2AgCI.3NH3 

Freezing 

Hot  Water    x  ^~^''    '  -^       ^    ^        ^    Mixture 


FIG.  33. 


ture,  hence  the  gas  may  be  liquefied  by  pressure.  A 
simple  apparatus  for  condensing  ammonia  is  shown  in 
Fig.  33. 

A  small  amount  of  the  compound  which  silver  chloride 
forms  with  ammonia,  viz.,  (AgCl)2.  3  NH3,  is  placed  in  the 
rounded  end  of  the  bent  tube.  The  tube  is  then  sealed,  and 
the  end  containing  the  silver  ammonium  compound  is  warmed 
in  a  water  bath,  while  the  drawn-out  end  is  cooled  in  a  freezing 
mixture.  The  compound  containing  the  ammonia  is  thus  de- 
composed, and  the  ammonia  is  condensed  in  the  drawn-out  end. 

Liquid  ammonia  is  about  0.6  as  heavy  as  water.  It 
boils  at  — 40°  C.  under  ordinary  pressure,  and  freezes 
at  —75°  C. 

146.  Liquid  Ammonia  as  a  Refrigerating  Agent.  - 

The  chief  use  of  liquid  ammonia  is  as  a  refrigerating 


LIQUID  AMMONIA  A  REFEIGEEATING  AGENT.      141 

agent,  and  to  produce  artificial  ice.  For  this  purpose 
gaseous  ammonia  (produced  by  heating  ammonia  water) 
is  condensed  to  the  liquid  state,  and  the  liquid  ammo- 
nia is  then  allowed  to  evaporate  rapidly.  The  heat 
necessary  for  the  vaporization  of  the  ammonia  is  absorbed 
from  the  body  to  be  cooled. 


Water  Supply 


MECHANICAL  COMPRESSION  SYSTEM. 
FIG.  34. 

The  apparatus  shown  in  Fig.  34  gives,  in  principle, 
ihe  construction  of  a  refrigerating  system. 

The  compression  engine  forces  gaseous  ammonia  into  con- 
densing pipes  under  so  great  a  pressure  that  the  ammonia  is 
liquefied.  The  condensing  pipes  are  cooled  by  streams  of 
water.  The  liquid  ammonia  is  then  allowed  to  expand  in  an- 
other system  of  pipes  surrounded  by  brine  (concentrated  salt 
solution)  ;  as  a  result  the  brine  is  ?ooled  to  a  little  above  its 


142  AMMONIA. 

freezing    point,   which   is  — 21°  C.     When    tanks    containing 
water  are  placed  in  the  cold  salt  solution,  the  water  is  frozen. 

Instead  of  ammonia,  liquid  sulphur  dioxide,  liquid 
carbon  dioxide,  compressed  air,  etc.,  are  often  used  to 
produce  artificial  ice. 

147.  Chemical    Properties.  —  Ammonia    does    not 
burn  in  the  air,  but  it  burns  readily  in  oxygen.     The 
products  are   chiefly  nitrogen  and  water,  according  to 
the  equation, 

2  NH3  +  3  O »  3  H2O  +  N2. 

Chlorine  reacts  energetically  with  ammonia,  forming 
nitrogen  and  hydrochloric  acid.  The  equation  is, 

2  NH3  +  3  C12 »  N2  +  6  HC1. 

If  the  ammonia  is  present  in  excess,  the  hydrochloric  acid 
unites  with  some  of  it,  forming  ammonium  chloride.  The  com- 
plete equation  is,  therefore, 

8  NH3  -f  3  C12 »  6  NH4C1  -f  N2. 

A  similar  reaction  occurs  between  ammonia  and  bromine.    . 

Ammonia  gas  and  hydrochloric  acid  gas  unite  in  equal 
proportions  by  volume  to  form  solid  ammonium  chloride. 

148.  Ammonium   Compounds.  —  In  the  compounds 
which   ammonia   forms    with  hydrochloric   acid,    nitric 
acid,  sulphuric  acid,  and  water,  the  ammonia  and  the 
other  substance  are  united  directly.     The  resulting  sub- 


A  MM  ONI  UM  COMPO  UNDS. 


143 


stances,  ammonium  chloride  (NH4C1),  ammonium  nitrate 
(NH4NO3),  ammonium  sulphate  [(NH4)2SO4],  and 
ammonium  hydroxide  (NH4OH),  might  be  called  "  am- 
monia hydrochlorate,  hydronitrate,"  etc.,  but  are  not. 
Instead,  they  are  looked  upon  as  compounds  of  the 
group  of  elements,  NH4,  with  the  element  Cl  and  the 
radicals  NO3,  SO4,  and  OH.  They  are  therefore  called, 
as  above,  ammonium  chloride,  nitrate,  etc. 

The  group  NH4  has  not  been  isolated,  since  it  breaks 
up  at  once,  on  being  liberated,  into  ammonia  and  hydro- 
gen ;  but  its  compounds  are  so  like  those  of  many  metals 
that  the  group  is  called  a  metallic  radical. 

To  what  extent  ammonium  resembles  the  metals  so- 
dium and  potassium  will  be  seen  by  a  comparison  of  some 
of  its  compounds  with  the  corresponding  compounds  of 
these  two  metals.  The  formulas  of  some  of  these  com- 
pounds are  given  below. 


AMMONIUM. 

SODIUM. 

POTASSIUM. 

Chloride. 

NH4C1 

NaCl 

KC1 

Nitrate. 

NH4N05 

NaNO;J 

KN03 

Sulphate. 

(NH4)2S04 

Na2S04 

K2S04 

Hydroxide. 

NH4OH 

NaOH                         KOH 

Carbonate. 

(NH4)2C03 

Na8C03 

K2C03 

Phosphate. 

(NH4),P04 

Na3P04 

K3P04 

144  AMMONIA. 

149.  Dissociation  of  Ammonium  Compounds.  —  All 

ammonium  compounds  decompose  readily,  liberating 
ammonia,  but  the  temperature  of  dissociation  varies 
greatly  in  different  cases. 

Thus,  the  union  of  ammonia  with  water  in  ammonium  hy- 
droxide is  so  weak  that  ammonium  hydroxide  cannot  be  iso- 
lated, although  it  is  probably  present  in  an  aqueous  solution  of 
ammonia. 

The  ammonium  salts  are  more  stable  than  the  hydrox- 
ide, but  even  these  decompose  readily. 

Ammonium  chloride,  for  example,  dissociates  completely  at 
350°  C.  into  ammonia  and  hydrochloric  acid  ;  above  350°  C., 
therefore,  ammonium  chloride  cannot  exist. 

When  the  substance  with  which  the  ammonia  is  com- 
bined is  not  volatile  at  the  temperature  of  dissociation, 
the  heating  of  an  ammonium  compound  simply  causes 
ammonia  to  be  libe rated. - 

Thus,  ammonium  phosphate  breaks  up  when  heated  into 
ammonia  and  phosphoric  acid.  Here  the  ammonia,  being  gase- 
ous, escapes  ;  but  the  non-volatile  phosphoric  acid  remains 
behind. 

If,  however,  the  substance  with  which  the  ammonia 
is  united  is  volatile  at  the  temperature  of  dissociation, 
the  ammonium  salt  sublimes. 

By  sublimation  we  mean  the  distillation  of  a  solid. 

This  is  the  case  with  ammonium  chloride  ;  heat  breaks  it  up 
into  ammonia  and  hydrochloric  acid,  but  both  these  sub- 
stances being  volatile,  they  pass  off  together  and  unite  again 
When  cold  to  form  solid  ammonium  chloride. 


SYNTHESIS    OF  AMMONIA.  145 

This  dissociation  of  ammonium  compounds  explains 
the  liberation  of  ammonia  from  ammonium  salts  by 
"  strong "  bases.  When  ammonium  chloride,  for  ex- 
ample, and  slaked. lime,  Ca(OH)2,  are  heated  together, 
the  ammonium  chloride  is  dissociated  into  ammonia  and 
hydrochloric  acid,  just  as  when  it  is  heated  by  itself. 
The  slaked  lime,  however,  "fixes  "  the  acid  by  forming 
with  it  calcium  chloride  and  water  (cf.  §  142).  As  a 
result,  only  ammonia  and  some  water  pass  off,  calcium 
chloride  remaining  behind. 

Similar  reactions  occur  when  sodium  hydroxide,  potassium 
hydroxide,  etc.,  are  used  in  place  of  slaked  lime. 

150.  Composition   of   Ammonia.  —  Ammonia     con- 
sists of  nitrogen  and  hydrogen  united  in  the  proportion, 
by  weight,  of  14  parts  nitrogen  to  3  parts  hydrogen. 
This  fact  is  indicated  by  the  formula  NH3. 

The  .volumetric  composition  of  ammonia  may  be 
proved  in  several  ways  :  — 

(1)  By  synthesis  from  nitrogen  and  hydrogen. 

(2)  By  the  action  of  chlorine  on  ammonia. 

151.  Synthesis  of    Ammonia  from  Nitrogen  and 
Hydrogen."  —  If  the  two  gases,  nitrogen  and  hydrogen, 
mixed  in  the  proportion  of  1  volume  of  nitrogen  to  3 
of  hydrogen,  are  subjected   to    the   electric    discharge, 
they  combine  in  part  to  form  ammonia  ;  the  union  can- 
not, however,  be  made  complete,  no  matter  how  long  the 
"sparking"  is  continued.      This  is  due  to  the  fact  that 
a  point  is  soon  reached  at  which  as  much  of  the  am- 


146  AMMONIA. 

monia  is  decomposed  in  a  given  time  as  is  formed  in 
the  same  time  by  the  union  of  nitrogen  and  hydrogen. 
To  make  the  combination  complete  it  is  necessary  to 
remove  the  ammonia  as  fast  as  it  is  formed. 

This  may  be  done  by  "  sparking  "  the  mixture  of  nitrogen 
arid  hydrogen  over  some  substance,  e.  (/%.,  sulphuric  acid,  capa- 
ble of  absorbing  the  ammonia.  Under  these  conditions  all  of 
the  nitrogen  and  hydrogen  disappears. 

It  is  thus  proved  that  1  volume  of  nitrogen  unites 
with  3  volumes  of  hydrogen  to  form 
ammonia.  The  union  is  accompanied 
by  a  shrinkage  in  volume,  4  volumes  of 
the  mixed  gases  becoming  2  volumes 
of  ammonia. 


152.  Action  of  Chlorine  on  Am- 
monia. —  Since  1  volume  of  chlorine 
unites  with  1  of  hydrogen  (cf.  §  94), 
the  relation  of  nitrogen  to  hydrogen 
can  be  obtained  readily  if  that  of  nitro- 
gen to  chlorine  can  be  established.  For 
this  purpose  we  make  use  of  the  known 
action  of  chlorine  upon  ammonia  (cf. 
§  147).  The  experiment  is  carried  out 
as  follows :  — 


FIG.  35. 


A  Hoffmann  tube  (Fig.  35)  is  filled  first 
with  a  saturated  solution  of  sodium  chloride  in  water,  and  then, 
by  displacement  of  the  salt  solution,  with  chlorine. 

The  stopcock  is  then  closed,  and  the  cup  above  the  stop- 


EXERCISES.  147 

cock  is  filled  with  concentrated  ammonium  hydroxide  solution. 
The  latter  is  now  run,  drop  by  drop,  into  the  chlorine,  care 
being  taken  that  no  air  enters  and  no  chlorine  escapes.  A 
flash  of  light  will  be  seen  when  the  first  drops  of  ammonia 
water  are  introduced. 

When  almost  all  of  the  ammonia  water  has  been  run  in, 
the  stopcock  is  closed  and  the  cup  is  filled  with  water.  The 
cup  is  now  connected,  by  means  of  a  tube  filled  with  water, 
with  a  beaker  of  dilute  sulphuric  acid,  the  stopcock  is  again 
opened,  and  the  water  and  dilute  acid  are  allowed  to  enter  un- 
til the  pressures  inside  and  outside  the  tube  are  equalized. 

The  tube,  which  was  full  of  chlorine,  will  now  be  one-third  full 
of  nitrogen.  Hence ,  — 

Volume  of  nitrogen  :  volume  of  chlorine  : :  1:3;  whence 

Yolume  of  nitrogen  :  volume  of  hydrogen  ::  1  :  3. 


153.  Exercises. 

1 .  How  many  grams  of  ammonia  gas  can  be  made  by  heat- 
ing 100  grams  ammonium  chloride  with  slaked   lime  ?     How 
many  liters  at  0°  C.  and  760  mm.  ? 

2.  What  wreight  of  ammonium  sulphate  is  required  to  give, 
with  lime,  20  liters  of  ammonia  gas  when  one  liter  of  ammonia 
weighs  0.77  grams  ? 

3.  What  weight  of  slaked  lime,  Ca(OH)2,  is    necessary  to 
decompose  50  grams  ammonium  nitrate  ?     How  many  grams 
ammonia  will  be  formed  ? 

4.  May  concentrated  sulphuric  acid  be  used  to  dry  ammo- 
nia gas  ?     Why  ? 

5.  How  could  you  separate  a  mixture  of  the  gases  ammonia, 
oxygen,  and  nitrogen  so  as  to  get  the  proportionate  amounts  of 
each  in  the  mixture  ? 

6.  What  would  be  the  volume  of  each  of  the  resulting  gases 


148  AMMONIA. 

if  500  c.c.  of  ammonia  gas  were  completely  decomposed  into 
its  constituents  ? 

7.  How  many  cubic  centimeters  of  oxygen  will  be  required 
to  unite  exactly  with  the  hydrogen  produced  by  the  complete 
decomposition  of  100  c.c.  of  ammonia  gas  ? 


CHAPTER  XII. 
NITROGEN   ACIDS   AND   OXIDES. 

154.  Nitric  Acid.  —  Nitric  acid  is  one  of  the  most 
important  substances  known  to  Chemistry.     It  has  been 
in  use  since  the  time  of  the  early  alchemists,  but  its  true 
nature  was  not  understood  until  the  latter  part  of  the 
eighteenth  century. 

During  the  Middle  Ages  nitric  acid  was  made  by  the 
distillation  of  a  mixture  of  alum,  blue  vitriol,  and 
niter. 

Alum  is  potassium  aluminum  sulphate  plus  crystal-water 
[K2SO4,  A12(SO4)3,  24  H2O],  and  blue  vitriol  is  cupric  sulphate 
plus  crystal-water  (CuSO4.  5  H2O)  ;  by  the  dry  distillation  of 
these  substances  sulphuric  acid  was  set  free.  This  with  the 
niter  (KNO3)  gave  nitric  acid. 

155.  Commercial  Preparation.  —  At  the  present  time 
nitric  acid  is  made  on  a  large  scale  by  heating  sodium 
nitrate  with  concentrated  sulphuric  acid.     The  operation 
is   carried    out  in    large   iron   retorts;  and    the   vapors 
evolved  are  condensed  in  a  system  of  earthenware  jars. 
The  resulting  liquid  is  redistilled. 

In  this  way  there  is  obtained  an  acid  of  specific  gravity  1.4 ; 
its  boiling  point  is  120°  to  121°  C.    This,  the  "  commercial "  grade 
.of  nitric  acid,  is  only  68%,  by  weight,  nitric  acid.     The  remain- 
der is  water. 

149 


150 


NITROGEN  ACIDS  AND    OXIDES. 


156.  Laboratory  Method. — In  the  laboratory,  nitric 
acid  is  commonly  made  by  heating  potassium  nitrate 
with  concentrated  sulphuric  acid  in  a  retort  or  distilling 

bulb  connected  with  a 
condenser.  The  con- 
denser may  be  a  test  tube 
partly  immersed  in 
water,  as  in  Fig.  36. 
When  no  more  acid  dis- 
tills over,  the  flask  is 
FIG.  36.  allowed  to  cool;  the 

other   product   of    the 

reaction,  potassium  hydrogen  sulphate   (KHSO4),  then 
crystallizes  out  in  the  form  of  long,  white  needles. 
The  equation  for  the  reaction  is,  — 


H2S04 


KHS04 
136 


101  98  136  63 

For  the   reaction    between   sodium   nitrate   and   sul- 
phuric acid  the  equation  is,  — 


H2SO4 


XaHSO4  -f 


157.  Preparation  of  Nitric  Acid  Compared  with 
that  of  Hydrochloric  Acid.  —  The  preparation  of  nitric 
acid  resembles  that  of  hydrochloric  acid.  In  each  case 
the  acid  is  set  free  from  its  salts  by  sulphuric  acid,  not 
because  sulphuric  acid  is  stronger  than  the  acid  it  dis- 
places, but  because  hydrochloric  acid  and  nitric  acid  are 
volatile,  and  are,  therefore,  removed  from  the  "  field  of 


PROPERTIES   OF  NITRIC  ACID.  151 

action."     As  a  result,  the  reaction  goes  on- until  practi- 
cally complete  (cf.  §  151). 

In  each  of  these  cases  only  half  of  the  hydrogen  of  sulphuric 
acid  is  replaced  by  the  metallic  element.  If  the  temperature 
is  raised  considerably,  a  second  portion  of  sodium  nitrate  will 
react  with  the  sodium  hydrogen  sulphate  formed,  giving  normal 
sodium  sulphate  and  a  second  portion  of  nitric  acid,  according 
to  the  equation, 

NaN03  +  NaHS04  =  Na28O4  +  HNO8. 

In  the  treatment  of  sodium  chloride  with  sulphuric  acid  on 
a  commercial  scale,  this  second  reaction  is  actually  carried  out, 
because  the  compound  ]^a2SO4  is  wanted.  In  the  preparation 
of  nitric  acid,  however,  the  second  reaction  is  of  no  advantage, 
since  the  high  temperature  necessary  decomposes  the  nitric 
acid. 

158.  Properties  of  Nitric  Acid. — Commercial  nitric 
acid  is  dehydrated  by  distilling  it  with  concentrated  sul- 
phuric acid.  The  anhydrous  acid  (the  best  yet  made  was 
probably  99.8^)  pure)  is  a  thick,  colorless  oil  of  specific 
gravity  1.56.  It  begins  to  distill  at  86°  C.,  but  breaks 
up  to  a  certain  extent  into  other  compounds.  The 
equation  for  its  decomposition  is,  — 

2  HN03 »  2  N02  (or  N2O4)  +  H2O  +  O. 

Nitrogen  dioxide  is  a  brown  gas  which  dissolves  readily  in 
nitric  acid  ;  hence  anhydrous  nitric  acid  that  has  been  distilled 
has  a  brown  color.  This  color  may  be  removed  by  bubbling 
air  through  the  liquid. 

The  same  decomposition  of  nitric  acid  takes  place  in  the 
light  (very  rapidly  in  sunlight)  ;  hence  the  nitric  acid  of  reagent 
bottles  soon  becomes  brown. 


152  NITROGEN  ACIDS  AND    OXIDES. 

When  the  vapors  of  nitric  acid  are  passed  through  a 
hot  tube  they  are  completely  decomposed  into  nitrogen 
dioxide,  water,  and  oxygen. 

As  an  acid,  nitric  acid  neutralizes  solutions  of  bases, 
forming  with  them  nitrates  am\  water,  just  as  hydrochloric 
acid  forms  chlorides-  and  water. 

Thus,  if  potassium  hydroxide  solution  is  treated  with  dilute 
nitric  acid  until  neutral,  and  the  solution  is  evaporated,  potas- 
sium nitrate  will  be  obtained.     The  equation  is,  — 
KOH  +  HNO3  =  KNO3  +  H2O. 


159.  Action  of  Nitric    Acid    upon  Metals.  —  The 

tendency  of  nitric  acid  to  break  up  into  oxides  of  nitro- 
gen and  free  oxygen  (<?/'.  §  158)  determines  its  general 
chemical  behavior  ;  for  nitric  acid  is  not  only  a  strong 
acid,  but  a  powerful  oxidizing  agent.  When,  therefore, 
we  compare  the  action  of  nitric  acid  upon  metals  with 
that  of  hydrochloric  and  dilute  sulphuric  acids,  we  ob- 
serve a  great  difference  ;  for  while  the  acids  just  named 
generally  give  up  hydrogen  when  treated  with  metals, 
nitric  acid  rarely  does  so.  Instead  of  being  set  free, 
the  nascent  hydrogen  usually  reduces  some  of  the  nitric 
acid  to  nitrogen  oxides  or,  even  lower,  to  hydroxylamine, 
NH2OH,  and  to  ammonia. 

Thus,  zinc  and  dilute  nitric  acid  (5%  to  6%)  give,  perhaps, 
zinc  nitrate  and  hydrogen,  according  to  the  equation, 

Zn  +  2  HN08  -  »  Zn(N03)2  +  H2  (c/.  §  9)  ; 
but  no  hydrogen  is  liberated.     Instead,  the  hydrogen  formed 
reduces  some  of  the  nitric  acid  to  ammonia.     Hence  the  solu- 


ACTION    OF  NITRIC  ACID    UPON  METALS.        153 

tion  contains,  besides  water,  zinc  nitrate  and  ammonium  nitrate. 
The  following  equations  probably  represent  what  takes  place  : 

(1)  4  Zn  +  8  HN03  =  4  Zn(NO3)2  +  4  H2 ; 
(This  is  the  equation  above  multiplied  by  4.) 

(2)  4  H2  +  HN03  =  NH4OH  +  2  H2O  ; 

(3)  NH^OH  -f  HNO3  =  NH4NO3  -f  H2O. 
Hence  the  complete  equation  must  be, 

4  Zn  +  10  HNO8  =  4  Zn(NO8)2  -f  NH4NO8  +  3  H2O. 

In  solutions  somewhat  more  concentrated  than  the 
above  the  reduction  of  nitric  acid  goes  only  to  nitrogen 
oxides.  The  equation  for  the  most  common  reduction 
of  nitric  acid  by  nascent  hydrogen  is,  probably,  — 

2  HN03  +  3  H2 >  4  H2O  +  2  NO. 

Copper  and  nitric  acid  of  specific  gravity  1.2  do  not 
react  according  to  the  equation, 

Cu  -f-  2  HN03  =  Cu(NO,)2  +  H2, 
for  the  reaction  is  an  oxidation :  — 

3  Cu  +  2  HN03 »  3  CuO  +  2  NO  +  H2O. 

Copper  nitrate  is  indeed  formed,  63  grams  of  copper 
giving  124  grams  of  copper  nitrate,  but  no  hydrogen  is 
evolved.  Instead  of  hydrogen  we  get  a  gas  which, 
though  colorless  itself,  forms  brown  fumes  when  it 
comes  in  contact  with  oxygen.  This  gas  is  nitric  oxide, 
NO,  and  the  brown  fumes  consist  of  nitrogen  dioxide, 
NO2 ;  hence  the  copper  must  have  reduced  some  of  the 
nitric  acid. 

The  cupric  oxide  then  reacts  with  the  nitric  acid,  giving 
cupric  nitrate  and  water. 

3  CuO  +  6  HNO3 »  3  Cu(NO3)2  +  3  H2O. 


154  NITROGEN  ACIDS  AND    OXIDES. 

The  complete  equation  is,  therefore, 

3  Cu  +  8  HNO3  =  3  Cu(NO3)2  +  4  H2O  +  2  NO. 
The  equation  for  the  action  of  silver  and  nitric  acid  is  simi- 
lar, viz. , — 

3  Ag  +  4  HNO3  =  3  AgNO3  -f  2  H2O  +  NO. 

160.  Aqua  Regia.  —  Gold  and  platinum  alone  of  all 
the  more  common  metals  do  not  react  with  nitric  acid. 
As    was    stated    under  Chlorine    (cf.  §  82),  these  two 
metals  are  soluble  in  aqua  regia  and  in  chlorine  water. 

Aqua  regia  is  made  by  mixing  nitric  acid  with  hydro- 
chloric acid  (three  volumes  of  hydrochloric  acid  to  one 
of  nitric  acid)  ;  it  is  merely  a  source  of  nascent  chlorine. 

Aqua  regia  is  used  to  dissolve  many  other  metals  be- 
sides gold  and  platinum. 

161.  Oxidation  by  Nitric  Acid.  — It  is  not  only  to- 
ward metals  that  nitric  acid  acts  as  an  oxidizing  agent. 
Phosphorus  is  converted  by  it  into  phosphoric  acid,  and 
sulphur  into   sulphuric    acid ;    while    glowing  charcoal 
burns  in  the  acid  much  as   it  would  in  oxygen  itself. 
Organic  coloring  matters,  e.  g.,  indigo  solution,  are  oxi- 
dized by  nitric  acid  to  colorless  bodies. 

The  oxidizing  power  of  nitric  acid  is  much  greater  if  some 
of  the  higher  oxides  of  nitrogen  are  present.  Such  an  acid  is 
fuming  nitric  acid,  a  red  liquid  containing  much  nitrogen  di- 
oxide, NO2.  This  acid  is  used  when  very  rapid  oxidation  is 
desired. 

162.  Explanation  of  Oxidation  by  Nitric  Acid.  - 

The  student  will  be  able  to  understand   the   peculiar 


FOEMATION  OF  NITRATES  IN  NATURE.        165 

behavior  of  nitric  acid  better  if  he  will  consider  nitric 
acid  as  made  up  of  water  and  nitrogen  pentoxide  ,  N9O., 
as  represented  by  the  equation, 


He  will  then  see  that  the  oxidizing  action  of  nitric  acid 
is  due  to  the  loss  of  oxygen  by  nitrogen  pentoxide.  Even 
without  a  reducing  agent,  nitrogen  pentoxide  readily 
gives  nitrogen  tetroxide  and  oxygen  (cf.  §  158)  ;  but 
in  the  presence  of  such  an  agent  the  decomposition  of 
nitrogen  pentoxide  is  very  easy  indeed.  Nitrogen  tet- 
roxide, N2O4  ;  nitrogen  trioxide,  N2O3  ;  nitric  oxide, 
NO  ;  nitrous  oxide,  N2O  ;  nitrogen,  N2,  or  even  ammo- 
nia, NH3,  may  be  formed,  according  to  circumstances. 

The   following   equations  represent  some  of  the   reactions 
that  may  take  place  :  — 
N2O5  =  N2O4  -f-  O.        This  takes  place  when  concentrated  nitric 

acid  is  heated,  or  treated  with  metals. 
N2O5  =  N2O3  -j-  2  O.    This  takes  place   with   starch,  arsenic 

trioxide,  etc.,  and  acid  of  specific  grav- 

ity 1.3. 
N2O5  =  2  NO  -f  3  O.    This    takes   place  with   copper,  silver, 

etc.,  and  acid  of  specific  gravity  1.2. 
NgOg  =  N2O  -|-  4  O.      This  takes  place  with  zinc  and  acid  of 

specific  gravity  1.1. 

163.  Formation  of  Nitrates  in  Nature.  —  We  have 
already  learned  (c/1  §  112)  that  nitrogen,  oxygen,  and 
water  vapor  are  combined  by  the  action  of  the  electric 
spark  into  nitric  acid  ;  consequently,  it  is  very  probable 
that  in  some  regions  much  nitric  acid  is  formed  in  this 


* 

156  NITROGEN  ACIDS  AND    OXIDES. 

way  and  gets  into  the  soil.  All  the  nitric  acid  of  com- 
merce, however,  is  made  from  the  nitrates;  these  are 
found  as  natural  deposits  in  certain  places. 

The  "  nitrate  beds "  were  probably  formed  by  the 
oxidation,  through  the  agency  of  bacilli,  of  nitrogenous 
organic  matter  in  the  presence  of  alkali.  The  most  ex- 
tensive deposits  of  nitrates  in  the  world  are  those  of  the 
Atacama  Desert  in  Chile.  The  alkali  present  is  sodium 
carbonate,  hence  Chile  saltpeter  is  sodium  nitrate  and 
not  the  potassium  salt. 

164.  Manufacture  of  Potassium  Nitrate.  —  Potas- 
sium nitrate  was  formerly  obtained  almost  exclusively 
from  Asiatic  countries,  where  it  appeared  as  a  deposit 
on  the  ground;  but  nowadays  most  of  it  is  made  from 
sodium  nitrate.  The  process  may  be  carried  out  as 
follows :  — 

Hot,  fairly  concentrated  solutions  of  potassium  chloride  and 
sodium  nitrate  are  mixed,  and  the  resulting  solution  is  poured 
off  from  the  crystalline  deposit  of  sodium  chloride  which  sepa- 
rates out.  The  solution  contains  much  potassium  nitrate  and 
small  amounts  of  sodium  chloride,  etc.  This  impure  potassium 
nitrate  is  then  redissolved  in  as  small  an  amount  of  hot  water 
as  possible,  and  allowed  to  crystallize  out ;  by  several  recrystal- 
lizations  pure  potassium  nitrate  is  obtained. 

The  equation  representing  the  formation  of  potassium  ni- 
trate is, — 

KC1  -f  NaKOg >  KKO3  +  KaCl. 

It  applies  only  to  a  mixture  of  concentrated  solutions  of  the 
factors. 


USES   OF  NITRIC  ACID  AND    THE  NITRATES.      157 

In  certain  European  countries  the  farmers  cultivate 
niter  by  introducing  the  proper  micro-organisms  into  a 
mixture  of  alkali  and  nitrogenous  matter. 

165.  Uses   of  Nitric   Acid    and    the    Nitrates.  - 

Nitric  acid  has  many  important  applications :  nitro- 
benzene, glyceryl  nitrate  (nitroglycerine),  and  nitrates 
of  cellulose  (collodion,  gun  cotton,  celluloid,  etc.)  are 
made  by  its  agency. 

The  nitrates  are  used  in  the  manufacture  of  gun- 
powder and  various  explosives,  and  in  the  preservation 
of  meats. 

Glyceryl  nitrate  (wrongly  called  nitroglycerine)  is  made  by 
the  action  of  a  mixture  of  concentrated  nitric  and  sulphuric 
acids  at  a  low  temperature  upon  glycerine.  It  is  a  thick, 
greenish  oil  of  a  very  unstable  nature,  and  very  explosive.  It 
is  usually  mixed  with  a  porous  earth,  and  appears  in  the  mar- 
ket chiefly  as  dynamite. 

Gun  cotton  and  collodion  are  made  by  the  action  of  a  mixture 
of  nitric  and  sulphuric  acids  upon  cotton. 

Celluloid  is  a  mixture  of  gun  cotton  and  camphor. 

Nitrobenzene  is  made  from  benzene,  C6H6,  and  the  nitric- 
sulphuric  acid  mixture.  When  nitrobenzene  is  reduced  by 
nascent  hydrogen,  it  gives  aniline,  the  starting  material  in  the 
manufacture  of  the  aniline  dyes. 

Potassium  nitrate  is  important  chiefly  as  a  constituent 
of  gunpowder,  which,  as  was  stated  in  §  26,  is  a  physi- 
cal mixture  of  potassium  nitrate  with  sulphur  and 
charcoal  in  various  proportions.  The  exact  amount  of 
each  of  the  constituents  depends  upon  the  purpose  for 


158  NITROGEN  ACIDS  AND    OXIDES. 

which  the  powder  is  to  be  used,  but  the  function  of  the 
potassium  nitrate  is  always  the  same,  viz.,  to  furnish 
oxygen  for  the  combustion  of  the  sulphur  and  the 
charcoal.  When  gunpowder  is  ignited,  it  forms  gases 
that  occupy  at  ordinary  pressure  several  hundred  times 
the  volume  of  the  original  gunpowder. 

The  reactions  taking  place  in  the  explosion  of  one  kind  of 
gunpowder  are  probably  represented  by  the  equation, 

2  KNO3  +  C  +  S »  K2SO4  -f  CO2  -f  N2. 

The  chief  use  of  potassium  nitrate  as  a  preservative 
is  in  the  preparation  of  "  corned  "  beef. 

166.  Nitrogen      Pentoxide.  —  Nitrogen     pentoxide 
("  penta  "  =  five)  is  of  theoretical  importance  because 
it  is  the  anhydride  of  nitric  acid,  i.  e.,  it  is  nitric  acid 
minus   water.     The  relation  of    nitrogen  pentoxide   to 
nitric  acid  is  evident  from  the  equation, 

N205  +  H20 »  2  HN03. 

Nitrogen  pentoxide  is  a  white,  crystalline  solid,  made  by  the 
distillation  of  anhydrous  nitric  acid  with  phosphorus  pen- 
toxide. 

167.  Nitrous  Acid.  —  Nitrous  acid  is  probably  con- 
tained in  a  solution  of  nitrogen  trioxide  in  water ;  ni- 
trogen trioxide  is,  therefore,  called  nitrous    anhydride, 
just  as  nitrogen  pentoxide  is  nitric  anhydride.     This  re- 
lation is  shown  by  the  equation, 

N203  +  H20 »  2  HN02. 


NITROGEN  TE  I  OXIDE.  159 

The  meaning  of  the  ending  ons  as  applied  to  nitrous  acid 
has  already  been  given  (cf.  §  106). 

Nitrous  acid  itself  has  not  been  made  ;  but  its  salts 
(called  nitrites  to  distinguish  them  from  the  salts  of 
nitric  acid)  are  formed  when  a  solution  of  nitrogen 
trioxide  is  neutralized  by  bases. 

Thus,  potassium  hydroxide  and  a  solution  of  nitrogen  tri- 
oxide give  potassium  nitrite  and  water,  as  is  shown  by  the 
equation, 

2  KOH  +  1SLO,  -  »  2  KXCX  4-  II0O. 

I  J      o  ti      I  2 

If  we  consider  the  solution  of  nitrogen  trioxide  to  be  nitrous 
acid,  the  equation  becomes,  — 

KOII  -f  Iim)2  =  K^sTO2  -f  II2O. 


A  second  way  in  which  nitrites  are  formed  is  by  the 
abstraction  of  oxygen  from  nitrates. 

Thus,  potassium  nitrate  loses  one-third  of  its  oxygen  when 
heated  to  a  high  temperature.     The  equation  is,  — 


The  temperature  required  is  much  lower  if  the  nitrate  is 
mixed  with  lead  when  heated  ;  for  the  lead  unites  with  the 
liberated  oxygen  to  form  lead  oxide,  PbO. 

168.  Nitrogen  Trioxide  (N203).  —  As  was  stated  in    • 
§  167,  nitrogen   trioxide   is   the    anhydride    of    nitrous 
acid.     It  may  be  made  by  the  action  of  nitric  acid  of 
specific  gravity  1.3  upon  arsenic  trioxide  or  upon  starch. 

The  brown  fumes  called  the  trioxide  may  be  condensed  by  a 
freezing  mixture  to  a  blue  liquid,  which  is  the  real  trioxide. 
When,  however,  this  liquid  is  distilled,  the  vapors  are  not  ni- 


160  NITROGEN  ACIDS   AND    OXIDES. 

trogen  trioxide,  but  a  mixture  of  the  dioxide  (NO2)  with  nitric 
oxide. 

*i69.     Nitrogen  Dioxide  and  Nitrogen  Tetroxide 

(N02  and  N204).  —  Nitrogen  dioxide  lies  between  nitro- 
gen trioxide  and  nitrogen  pentoxide  as  regards  the  pro- 
portion of  oxygen  it  contains.  Below  22°  C.  it  is  a 
liquid ;  but  it  is  usually  known  in  the  form  of  its  vapor. 
Nitrogen  tetroxide  ("  tetra  "  —  four)  exists  at  low 
temperatures,  but  it  dissociates  readily,  to  some  extent 
even  at  0°  C.,  into  nitrogen  dioxide.  This  is  shown  in 
the  equation, 

N204  =  2  N02. 

The  degree  of  the  dissociation  is  shown  by  the  dark- 
ening of  the  color,  nitrogen  tetroxide  being  colorless, 
but  nitrogen  dioxide  brown. 

When  nitrogen  tetroxide  is  dissolved  in  much  cold  water, 
the  solution  contains  a  mixture  of  nitrous  arid  nitric  acids  ; 
hence  nitrogen  tetroxide  is  the  anhydride  of  both  of  these 
acids.  The  equation  showing  this  fact  is,  — 

N2O4  (=  2  NO2)  -f  H2O »  HNO2  +  HXO3. 

Nitrogen  tetroxide  is  formed  when  lead  nitrate, 
Pb(NO3)2,  is  heated.  The  vapors  may  be  condensed 
by  passing  them  through  a  U-tube  surrounded  by  a 
freezing  mixture.  The  decomposition  of  lead  nitrate 
corresponds  exactly  with  that  of  nitric  acid,  as  is  shown 
by  the  equation, 

Pb(N03)2  =  PbO  +  N204  +  O  (_of.  §  158). 

*The  names  used  are  in  accordance  with  the  latest  nomenclature. 


NITRIC   OXIDE.  *  161 

170.  Nitric  Oxide  (NO). — Two  other  oxides  of  ni- 
trogen are  known,  viz.,  nitric  oxide  and  nitrous  oxide ; 
.both  are  colorless  gases  at  the  ordinary  temperature  and 
pressure. 

Nitric  oxide  is  produced  when  nitric  acid  of  specific 
gravity  1.2  is  allowed  to  react  with  copper.  The  equa- 
tion has  already  been  given  (</.  §  159).  It  is,  — 

3  Cu  +  8  HN03  =  3  Cu(N03)2  +  4  H2O  -f  2  NO. 

Nitric  oxide  is  slightly  heavier  than  air.  One  liter 
of  it  weighs,  under  standard  conditions,  1.34  grams. 
The  gas  is  only  slightly  soluble  in  water. 

When  nitric  oxide  is  passed  into  a  solution  of  ferrous  sul- 
phate, FeSO4,  large  quantities  of  the  gas  are  absorbed,  a  brown 
compound  of  ferrous  sulphate  and  nitric  oxide  being  formed, 
When  the  solution  containing  this  compound  is  heated,  pure 
nitric  oxide  escapes. 

Nitric  oxide  becomes  brown  when  it  comes  into  con- 
tact with  air  or  oxygen,  owing  to  the  formation  of  nitro- 
gen dioxide. 

One  volume  of  nitric  oxide 
consists  of  one-half  a  volume 
of  nitrogen  and  one-half  a 
volume  of  oxygen. 

The  amount  of  the  nitrogen 

,       ,      , .  FIG.  37. 

may   be   shown    by   heating    a 

piece  of  metallic  sodium  in  a  measured  volume  of  the  gas  over 
mercury  (see  Fig.  37).  The  sodium  .combines  with  the  oxygen 
of  the  nitric  oxide,  leaving  the  nitrogen.  The  volume  of  the 
nitrogen  should  be  just  half  that  of  the  nitric  oxide  taken. 


162  NITROGEN  ACIDS  AND    OXIDES. 

Nitric  oxide  supports  the  combustion  of  phosphorus 
and  of  magnesium,  as  well  as  of  sodium. 

171.  Nitrous  Oxide  (N20). — Nitrous  oxide  is  com- 
monly made  by  heating  ammonium  nitrate  (NH4NO3) 
to  170°  C.,  or,  better,  by  heating  a  mixture  of  sodium 
nitrate  and  ammonium  sulphate  to  a  slightly  higher 
temperature. 

The  formation  of  nitrous  oxide  is  the  result  of  sev- 
eral reactions.  At  first  the  ammonium  nitrate  probably 
dissociates  into  ammonia  and  nitric  acid  just  as  ammo- 
nium chloride  gives  ammonia  and  hydrochloric  acid. 
The  ammonia  and  nitric  acid  do  not,  however,  recom- 
bine  when  cool,  because  the  nitric  acid  oxidizes  the 
hydrogen  of  the  ammonia  to  water. 

The  final  equation  for  the  decomposition  of  ammo- 
nium nitrate  is,  — 

NH4NO3 »  N2O  +  2  H2O. 

Compare  with  this  the  equation  for  the  decomposition  of 
ammonium  nitrite  (§  111).  Note  that  all  of  the  hydrogen  is 
oxidized  to  water  in  each  case,  and  that  it  is  the  oxygen  in  ex- 
cess of  that  required  to  oxidize  hydrogen  that  causes  the  for- 
mation of  nitrous  oxide  when  ammonium  nitrate  is  decomposed. 

Nitrous  oxide  is  the  only  gas  besides  oxygen  that  will 
re-ignite  a  glowing  splinter. 

The  volumetric  composition  of  nitrous  oxide  may  be 
determined  in  the  same  way  as  that  of  nitric  oxide, 
viz.,  by  taking  out  the  oxygen  by  means  of  sodium.  There 
is,  however,  a  great  difference  in  results  ;  for  while  one 


EXERCISES.  ]  63 

0 

volume  of  nitric  oxide  contains  one-half  a  volume  of 
nitrogen  and  one-half  a  volume  of  oxygen,  one  volume 
of  nitrous  oxide  contains  one  volume  of  nitrogen  and 
one-half  a  volume  of  oxygen. 

Nitrous  oxide  may  be  condensed  to  the  liquid  form  ;  as  such 
it  is  a  commercial  article.  When  the  pressure  under  which 
the  gas  is  kept  is  removed,  some  of  the  liquid  vaporizes,  pro- 
ducing the  "  laughing  gas  "  which  dentists  use  to  produce  in- 
sensibility to  pain. 

Nitrous  oxide  is  easily  soluble  ;  1  volume  of  water 
at  0°  C.  absorbs  1.3  volumes  of  the  gas.  One  liter  at 
standard  conditions  weighs  1.97  grams. 

172.  Hyponitrous     Acid,    (NOH)2.  —  Hyponitrites, 
i.  e.,  salts  of   hyponitrous  acid,  have   been   known   foi 
some  time,  but  the  acid  itself  has  only  recently  been 
isolated.     A   solution  of   hyponitrous  acid  decomposes 
readily,  giving  nitrous  oxide,  according  to  the  equation, 

(NOH)2 >N20  +  H./>. 

Nitrous  oxide  may  thus  be  looked  upon  as  in  some  sense 
the  anhydride  of  hyponitrous  acid ;  but  the  union  of  nitrous 
oxide  and  water  to  form  the  acid  does  not  take  place. 

173.  Exercises. 

1.  How  many  grams  of  nitric  acid  can  be  made,  theoreti- 
cally, from  1  kg.  sodium  nitrate  with  sulphuric  acid  ? 

2.  What  will  be  the  volume  of  28.4  grams  of  commercial 
nitric  acid  ? 


164  NITROGEN  ACIDS  AND    OXIDES. 

3.  How  much  sulphuric  acid  (calculated  as  100%  H28O4)  will 
be  needed  to  give  1,260  grams  of  nitric  acid  with  potassium 
nitrate,  if  potassium  hydrogen  sulphate  is  formed  ? 

4.  Calculate  the  percentage  composition  of  nitric  acid  ? 

5.  How  much  potassium  hydroxide  is  required  to  neutralize 
exactly  a  solution  containing  42  grams  nitric  acid  ? 

6.  How  could   you  separate  a  mixture   of  silver  and  gold 
chemically  ? 

7.  How  many  grams  nitrogen  tetroxide  could  be  made  from 
450  grams  lead  nitrate  ?     How  much  oxygen  ? 

8.  How  many  grams  nitric  oxide  can  be  made,  theoretically, 
from  100  grams  nitric  acid  with  copper?     How  much  copper  is 
needed?     How  much  cupric  nitrate  is  formed? 

9.  How  many  liters  of  nitrous  oxide  at  0°  C.  and  760  mm. 
can  be  made  from  240  grams  ammonium  nitrate  ?     How  many 
at  20°  C.  and  720  mm.  ? 

10.  How  would  you  distinguish  between  nitrous  oxide  and 
oxygen  ? 

n.  How  many  cubic  centimeters  of  each  of  its  constituents 
combine  to  form  100  c.c.  nitric  oxide  ?  To  form  100  c.c.  of 
nitrous  oxide  ? 

12.  Name  three  different  classes  of  nitrates,  basing  the  dif- 
ference upon  the  way  in  which  the  members  of  each  class 
decompose  when  heated. 


CHAPTER   XIII. 


SULPHUR  AND  ITS  COMPOUNDS. 

174.  Occurrence  and  Preparation  of  Sulphur.  —  SuL 

phur  occurs  in  nature  in  both  a  free  and  a  combined 
form.  In  the  free  condition  it  is  obtained  chiefly  from 
Sicily,  Mexico,  and,  to  some  extent,  from  Louisiana. 
Natural  sulphur  is  usually  found  mixed  with  much 
earthy  material,  from  which  it  must  be  separated  to  pre- 
pare it  for  the  market. 

The  first  operation  in  the 
purification  of  sulphur  usu- 
ally consists  in  heating  the 
natural  product ;  the  sul- 
phur melts  and  flows  away, 
leaving  the  infusible  im- 
purities behind. 

In  the  second  operation 
the  partially  purified  sul- 
phur is  distilled  from  large 
iron  retorts  (see  Fig.  38), 
and  is  thus  separated  from  FIG.  gg. 

less  volatile  impurities. 

The  melted  sulphur  in  the  reservoir  A  is  allowed  to  flow 
from  time  to  time  into  the  retort  5,  in  which  the  sulphur  is 
vaporized.  The  sulphur  vapor  which  passes  into  the  con- 

165 


166  SULPHUR  AND   ITS   COMPOUNDS. 

denser  collects  either  in  the  liquid  state,  at  the  bottom  ((7)  of 
the  condenser,  or  in  a  solid  state  upon  the  cold  walls  (D).  The 
liquid  sulphur  is  run  into  molds  to  crystallize,  thus  producing 
the  "  roll-sulphur,"  or  "  brimstone  "  of  commerce  ;  the  sul- 
phur which  solidifies  upon  the  walls  appears  in  the  form  of  fine 
meal  and  is  called  "flowers "  (more  correctly,  "flour")  of 
sulphur. 

175.  Physical    Properties. — Sulphur,    like    many 
other    elements,    exists    in    several    different   physical 
forms;    consequently,   in  giving  the  properties   of  sul- 
phur we  must  specify  the  kind  of  sulphur  to  which  we 
are  referring.     The  several  varie- 
ties of    sulphur    may  he  grouped 
into  three  classes  :  — 

(1)  Ordinary,  or   rhombic,  sul- 
phur (Fig.  39).     This  is  the  form 
that  occurs  in   nature.     All  other 
forms  revert  to  this  form.     Its  spe- 
FIG'  39>  cine  gravity  is  2.07. 

(2)  Prismatic    sulphur    (Fig.    40). 
This  is  formed  by  the  slow  cooling  of 
fused  sulphur  of  any  of  the  other  varie- 
ties.    Its  specific  gravity  is  1.96. 

(3)  Amorphous  sulphur.     This  form 
is  produced  when    sulphur  at   tempera- 
tures above   230°  C.  is  chilled  rapidly,  as  by  pouring  it 
into  cold  water. 

At    ordinary    temperatures     amorphous    sulphur    changes 
slowly  (after  some  days)   into  the   ordinary  form ;   at   about 


CHEMICAL   PROPERTIES.  167 

100°  C.,  however,  the  change  is  instantaneous,  and  much  heat 
is  evolved.  Ordinary  and  prismatic  sulphur  are  readily  soluble 
in  carbon  disulphide,  CS2 ;  but  amorphous  sulphur  is  partly  in- 
soluble in  that  liquid. 

The  form  known  as  "  flowers"  of  sulphur  is  both  crystalline 
and  amorphous.  It  consists  of  a  crystalline  kernel  and  an 
amorphous  covering. 

The  existence  of  an  element  in  several  forms  is  called 
allotropism,  and  the  different  varieties  of  the  element 
are  called  its  allotropic  forms  (<?/*.  §  263). 

Ordinary  sulphur  has  a  yellow  color  and  is  practically 
without  odor  and  taste.  It  is  soluble  to  a  slight  extent 
in  ether  and  in  alcohol,  but  is  insoluble  in  water.  The 
best  solvent  for  sulphur  is  carbon  disulphide;  100  parts 
of  this  substance  dissolve  46  parts  of  sulphur  at  the 
ordinary  temperature. 

Ordinary  sulphur  behaves  peculiarly  when  heated.  At 
114°C.  it  melts,  becoming  a  yellow  liquid.  As  the  heating  is 
continued  the  sulphur  becomes  more  viscous  and  dark-colored  ; 
at  250°  C.  it  is  almost  black  and  can  hardly  be  poured.  Above 
300°  C.  it  is  limpid  again,  and  at  448°  it  boils,  forming  a  yellow 
vapor. 

176.  Chemical  Properties.  —  Sulphur  unites  directly 
with  many  elements,  especially  with  metals.  Thus, 
when  a  mixture  of  powdered  iron  and  sulphur  is  heated, 
or  moistened  with  water,  or  strongly  compressed,  it 
unites  chemically,  forming  ferrous  sulphide,  FeS.  Simi- 
larly, copper  foil  burns  in  sulphur  vapor,  giving  copper 
sulphide.  Mercury  combines  with  sulphur  when  the 


168  SULPHUR  AND  ITS   COMPOUNDS. 

two  are  simply    triturated,  i.  e.,  rubbed  together  in  a 
mortar. 

There  are  many  common  illustrations  of  the  action  of  sulphur 
upon  metals.  Thus,  silver  egg-spoons  are  blackened  by  the 
sulphur  contained  in  eggs  ;  and  silver  coins  which  have  been 
carried  about  in  the  pockets  are  colored  by  the  sulphur  of  the 
perspiration. 

The  illuminating  gas  of  many  cities  contains  sulphur  com- 
pounds ;  this  fact  accounts  for  the  tarnishing  of  articles  of  brass, 
copper,  lead,  silver,  etc.,  in  houses  in  which  such  gas  is  used. 

At  about  260°  C.  sulphur  begins  to  unite  with  the 
oxygen  of  the  air.  If  sulphur  at  a  temperature  slightly 
below  260°  C.  is  examined  in  the  dark  it  will  be  found 
to  phosphoresce,  i.  e.,  glow. 

In  burning,  sulphur  produces  sulphur  dioxide,  SO9. 
This  is  an  invisible  gas  having  the  characteristic  odor 
of  burning  sulphur. 

177.  Uses  of  Sulphur.  —  Sulphur  is  used  in  the 
preparation  of  many  important  substances.  Thus,  rub- 
ber goods  and  vulcanite  are  made  by  heating  together 
caoutchouc  and  sulphur ;  match  tips,  especially  the  older 
forms,  contain  sulphur  as  an  ingredient ;  gunpowder,  as 
previously  stated,  is  a  mixture  of  charcoal,  sulphur,  and 
niter;  and, sulphur  dioxide,  which  is  used  for  bleaching 
and  as  a  germicide,  is  made  by  burning  sulphur  in  air. 

Finally,  sulphur  is  used  in  the  preparation  of  sul- 
phuric acid,  which  is  possibly  the  most  important  sub- 
stance manufactured. 


HYDROGEN  SULPHIDE.  169 

178.  Compounds  of  Sulphur.  —  Sulphur  is  a  constit- 
uent of  many  important  compounds  ;  for  the  sulphides, 
next  to  the  oxides,  are   the  most  common  ores   of  the 
metals.     Among  the  natural'  sulphides  are  iron  pyrites 
(FeS2),  galena    (PbS),  and   hydrogen  sulphide   (H2S). 
The  latter  is  a  gas  under  ordinary  conditions. 

The  two  important  oxides  of  sulphur  are  sulphur 
dioxide  and  sulphur  trioxide. 

With  hydrogen  and  oxygen  sulphur  forms  sulphuric 
acid  and  other  acids,  and  with  metals  and  oxygen,  sul- 
phates, sulphites,  thio  sulphates,  etc. 

The  most  important  natural  sulphate  is  gypsum, 
CaSO^.  2  H2O.  Natural  barium  sulphate,  BaSO4,  is 
called  "  heavy  spar"  Both  are  important  minerals. 

Iron  pyrites,  FeS2,  is  a  source  of  both  sulphur  and  sulphur 
dioxide.  When  it  is  roasted,  i.  e.,  heated  in  a  current  of  air, 
its  sulphur  is  oxidized  to  sulphur  dioxide,  SO2,  but  if  the  iron 
pyrites  is  heated  without  access  of  air  it  breaks  down  into  a 
compound  of  iron  and  sulphur  containing  only  two-thirds  of 
the  sulphur  of  the  original  pyrites.  The  excess  of  sulphur  is 
liberated. 

The  equation  representing  this  reaction  is, — 

3  FeS2  =  Fe3S4  +  2  S. 

The  reaction  is  like  that  which  takes  place  when  manganese 
dioxide  is  heated  (cf.  §21),  and  which  is  indicated  by  the 
equation, 

3MnO2=Mn3O4-fO2. 

179.  Hydrogen  Sulphide. — Hydrogen  sulphide,  or 
hydrosulphuric   acid,   is    a   colorless    gas    composed    of 


170  SULPHUR  AND  ITS   COMPOUNDS. 

hydrogen  and  sulphur  in  the  proportion  of  1  part  by 
weight  of  hydrogen  to  16  parts  of  sulphur.  This  fact 
is  shown  by  the  formula,  H2S. 

Hydrogen  sulphide  is  usually  prepared  by  the  action 
of  dilute  sulphuric  or  hydrochloric  acid  upon  iron  sul- 
phide. The  reactions  are  shown  by  the  equations, 

FeS  -f-  H2SO4 >  FeSO4  -f  H2S,  and 

FeS  +  2  HC1 »  FeCl2  -f  H2S. 

The  iron  sulphate  or  chloride  formed  remains  in  solu- 
tion. 

The  method  just  described  gives  hydrogen  sulphide  in  a 
form  good  enough  for  ordinary  use,  but  not  pure.  Pure  hydro- 
gen sulphide  is  prepared  by  the  action  of  concentrated  hydro- 
chloric acid  upon  antimony  trisulphide,  Sb2S3.  The  equation 
is,— 

Sb2S3  -f  6  HC1 »  2  SbCl8  +  3  H2S. 

1 80.  Properties  of  Hydrogen  Sulphide.  —  Hydrogen 
sulphide  has  the  odor  of  rotten  eggs.  It  is  formed  by 
the  decomposition  of  most  organic  substances  containing 
sulphur.  So-called  "sulphur"  waters  OAve  their  proper- 
ties to  the  hydrogen  sulphide  dissolved  in  them. 

The  gas  is  1.18  times  as  heavy  as  air. 

Hydrogen  sulphide  is  very  soluble  in  water;  one 
volume  of  water  absorbs  at  standard  conditions  three 
volumes  of  the  gas.  The  aqueous  solution  is  readily 
decomposed,  especially  in  the  light  and  when  warm,  by 
the  oxygen  of  the  air.  The  hydrogen  is  thus  converted 


SULPHIDES.  171 

into  water,  and  the  sulphur  is  set  free ;  consequently  a 
solution  of  hydrogen  sulphide  soon  loses  its  odor  and 
deposits  sulphur.  The  equation  is,  — 


2  H2S  +  (X  -  >  2  HoO  +  2  S. 

Similar  to  the  action  of  oxygen  is  that  of  chlorine,  which 
forms  with  hydrogen  sulphide  or  its  aqueous  solution  hydro- 
chloric acid  and  sulphur,  according  to  the  equation, 

II2$  +  Cl,  -  »  2  HC1  +  S. 
Iodine  acts  in  the  same  way,  viz.:  — 


This  method  is  used  for  the  preparation  of  hydriodic  acid. 
Hydrogen  sulphide  is,  as  we  might  expect,  a  reducing  agent. 

181.  Sulphides.  —  Hydrogen  sulphide  is  a  weak  acid, 
and  is  therefore  called  hydro  sulphuric  acid.  Its  salts,  the 
sulphides,  may  be  formed  in  several  ways  ;  these  are,  — 

(1)  By  the  reduction  of  a  sulphate  or  sulphite  ; 

(2)  By  the  neutralization  of  an  alkali  with  hydrogen 
sulphide  ; 

(3)  By  the  addition  of  hydrogen  sulphide  to  a  soluble 
salt  of  the  metal  whose  sulphide  is  to  be  formed. 

An  illustration  of  the  first  method  is  the  reduction  of 
sodium  sulphate,  when  heated  with  charcoal,  to  sodium 
sulphide.  The  equation  is,  — 


Na2S04  +  4  C  -  >  Na2S  +  4  CO. 
An  illustration  of  the  second  method  is  the  absorption 


172  -SULPHUB  AND  ITS   COMPOUNDS. 

of  hydrogen  sulphide  by  a  solution  of  sodium  hydroxide. 
The  equation  is,  — 

2  KaOH  -f  H2S >  Xa2S  +  2  H2O. 

Hydrosulphides. — The  equation  just  given  represents  only 
the  final  products  formed  in  the  ordinary  method  of 'preparing 
sodium  sulphide  ;  for,  if  we  pass  hydrogen  sulphide  into  sodium 
hydroxide  solution  until  no  more  hydrogen  sulphide  is  absorbed, 
we  obtain  sodium  hydrosulphide,N&SH.  The  equation  is, — 

KaOH  -f  H2S  =  NaSH  +  H2O. 

If  we  now  add  to  the  sodium  hydrosulphide  as  much  sodium 
hydroxide  as  we  used  originally,  we  shall  obtain  sodium  sul- 
phide, Na2S,  as  is  represented  in  the  equation, 

NaOH  -f  NaSH  =  !Na2S  +  H2O. 

To  prepare  ammonium  sulphide,  £N"II4)2S,  we  employ  a  simi- 
lar method. 

182.  Precipitation  of  Sulphides.  —  The  third  method 
of  forming  a  sulphide,  viz.,  by  adding  hydrogen  sul- 
phide to  a  soluble  salt  of  the  metal,  will  succeed  only  in 
case  the  sulphide  desired  is  insoluble  in  the  solvent  present. 
Thus  if  hydrogen  sulphide,  either  in  gaseous  form  or 
in  aqueous  solution,  is  added  to  a  solution  of  cupric  sul- 
phate, CuSO4,  cupric  sulphide  and  sulphuric  acid  are 
formed,  according  to  the  equation, 

CuSO4  +  H2S >  CuS  -f  H2SO4. 

The  cupric  sulphide  will  appear  as  a  black  precipitate. 
In  future  we  shall  usually  italicize  the  formula  of  a  precipitate, 


PRECIPITATION   OF  SULPHIDES.  173 

Explanation.  —  The  student  will  notice  that  in  every 
case  the  action  of  hydrogen  sulphide  upon  the  salt  of  a 
metal  must  tend  to  produce  a  free  acid  in  addition  to 
the  sulphide  of  the  metal.  Now,  we  know  that  free  acids 
generally  act  upon  sulphides  giving  hydrogen  sulphide, 
this  being  in  fact  the  way  in  which  hydrogen  sulphide 
is  prepared  (cf.  §  179) ;  hence  the  precipitation  of 
cupric  sulphide  from  a  solution  of  cupric  sulphate  by 
hydrogen  sulphide  must  be  possible  only  because  the 
cupric  sulphide  is  insoluble  in '  the  dilute  acid  formed  at 
the  same  time. 

Sulphides  that  are  soluble  in  dilute  acids  cannot,  therefore, 
be  precipitated,  or  only  incompletely,  by  hydrogen  sulphide. 

Thus,  no  precipitate  is  produced  when  hydrogen  sulphide  is 
added  to  manganese  sulphate  solution,  because  the  reverse  re- 
action, represented  by  the  equation, 

MnS  +  H2SO4 »  MnSO4  -f  H2S, 

is  the  one  that  tends  to  take  place.  If,  however,  we  use,  in- 
stead of  hydrogen  sulphide,  a  soluble  salt  of  hydrogen  sulphide, 
precipitation  of  manganese  sulphide  occurs,  for  the  reaction 
cannot  produce  free  acid. 

Thus,  with  sodium  sulphide  the  reaction  is  represented  by 
the  equation, 

MnSO4  -f  Na2S >  MnS  -f  Na2SO4. 

Some  sulphides,  however,  are  soluble  in  water  itself  ;  such 
sulphides  cannot,  of  course,  be  precipitated  by  either  hydrogen 
sulphide  or  its  salts.  Barium  and  calcium  sulphides  are  ex- 
amples. 


174  SULPHUR  AND  ITS   COMPOUNDS. 

183.  Carbon  Bisulphide  (CS2).  —  Carbon  disulphide 
is  formed  by  the  direct  union  of  carbon  and  sulphur, 
the  usual  method  being  to  pass  sulphur  vapor  over  hot 
charcoal. 

When  pure,  carbon  disulphide  is  colorless  and  has  an 
ethereal  odor ;  but  as  obtained  commercially  it  is  often 
yellow  and  has  a  disagreeable  smell.  The  liquid  boils 
at  47°  C. 

Carbon  disulphide  is  very  inflammable.  The  prod- 
ucts of  its  combustion  are  carbon  dioxide  and  sulphur 
dioxide,  as  represented  by  the  equation, 

CS2+302  =  C02  +  2S02. 

The  chief  use  of  carbon  disulphide  is  as  a  solvent  for  sul- 
phur, caoutchouc,  phosphorus,  iodine,  etc. 

184.  Manufacture  of  Sulphuric  Acid.  -  -  The  mod- 
ern method  of  making  sulphuric  acid  is  to  treat  sulphur 
trioxide  with  water  (cf.  §  190).     The    so-called  "  Kng- 
lish,"  or  common,  process  consists  in  oxidizing  sulphur 
dioxide   in  the  presence  of  water. 

The  oxidizing  agent  used  is  nitric  acid.  The  sulphur 
dioxide  is  produced  from  sulphur,  iron  pyrites  (FeS2), 
or  galena  (PbS);  its  oxidation  is  carried  out  in  large 
boxes  lined  with  lead  and  called  "  the  leaden  cham- 
bers." Currents  of  air,  of  steam,  and,  occasionally, 
of  nitric  acid  enter  the  leaden  chambers  along  with 
the  sulphur  dioxide,  and  sulphuric  acid  is  the  result. 
The  simplest  equation  is,  — 

S02  +  H2O  +  0 »  H3S04. 


MANUFACTURE   OF  SULPHURIC  ACID. 


175 


Explanation.  —  The  nitric  acid  introduced  ii>to  the  leaden 
chambers  is  reduced  to  nitric  oxide,  NO;  a  small  amount  of 
sulphur  dioxide  is  thus  oxidized  directly  by  nitric  acid.  But 
the  greater  portion  of  the  oxygen  used  comes  from  the  air ;  for 


FIG.  41. 


the  nitric  oxide  takes  up  the  oxygen  of  the  air,  forming  nitro- 
gen dioxide  (c/.  §  170),  and  then  gives  up  the  oxygen  to  the 
sulphur  dioxide.  All  of  these  facts  are  represented  by  the 
following  equations  :  — 


(1) 
(2) 
(3) 
(4)  Repetition  of  (2). 

Theoretically,  a  very  small  amount  of  nitric  acid 
ought  to  be  able  to  oxidize  an  indefinitely  large  amount 
of  sulphur  dioxide,  but  in  practice  some  of  the  nitrogen 
oxides  are  lost  ;  hence  nitric  acid  must  be  added  from 


176  SULPHUR  AND  ITS   COMPOUNDS. 

time  to  time  to  the  mixture  in  the  leaden  chambers. 
The  greater  portion  of  the  nitrogen  oxides  is  prevented 
from  escaping  by  being  made  to  pass  through  towers  of 
dilute  sulphuric  acid,  which  absorbs  them.  They  are 
thus  compelled  to  perform  the  work  of  oxidizing  sul- 
phur dioxide  over  and  over  again. 

The  equations  given  above  are  only  incomplete  representa- 
tions of  the  reactions  taking  place  in  the  leaden  chambers. 

It  is  probable  that  the  nitric  oxide  reacts  with  the  steam, 
oxygen  and  sulphur  dioxide  present  in  the  leaden  chambers, 
giving  a  substance  called  nitrosyl  sulphuric  acid.  The  equa- 
tion is,  — 

2  SO2  +  2  NO  +  3  O  -f  IT./)  =  -2  XO.  TISO4. 

The  nitrosyl  sulphuric  acid  is  a  solid  substance.  It  is  known 
technically  as  "  chamber  crystals."  It  is  readily  decomposed 
by  the  excess  of  steam,  giving  sulphuric  acid  and  nitrogen 
trioxide  (N"0O3),  as  is  shown  by  the  equation, 

2  NO.  HSO4  +  H2O  =  2  II2SO4  -f  N,O3- 
Apparatus  for  demonstrating    the  nitrosyl  sulphuric 
acid  manufacture  is  shown  in  Fig.  41. 

185.  Purification  of  Sulphuric  Acid. — The  sul- 
phuric acid  obtained  in  the  leaden  chambers  contains 
about  40^o  of  water;  it  is  therefore  concentrated 
by  evaporation.  The  evaporation  is  carried  out  in 
leaden  pans  until  the  acid  is  concentrated  enough  to 
attack  the  lead.  When  this  is  the  case,  further  evapora- 
tion is  carried  out  in  cast-iron  pans  until  an  acid  con- 
taining about  13J&  of  water  is  obtained.  This  is 


REDUCTION   OF  SULPHURIC  ACID.  177 

the  "crude"  sulphuric  acid  of  commerce.  It  is  very 
impure.  If  the  acid  is  to  be  concentrated  still  further, 
the  evaporation  must  be  carried  out  in  vessels  of  glass, 
porcelain,  or  platinum. 

The  pure  acid  is  made  by  distilling  the  crude  product  in 
stills  of  platinum  lined  with  gold.  There  is  thus  obtained  an 
oil  boiling  at  338°  C.  and  containing  only  1.5%  water.  Its 
specific  gravity  is  1.S54  at  0°  C.  The  anhydrous  acid 
(approximately  100%  sulphuric  acid)  is  made  only  in  very 
small  amounts. 

1 86.  Properties.  —  Sulphuric   acid    is  a  thick,    oily, 
colorless  liquid.     When  it  is  diluted  with  water,  much 
heat  is  evolved,  so  much,  indeed,  that  the  water  some- 
times boils.      To  avoid  spattering  of  the  hot  liquid  we 
pour  the  concentrated  acid  in  a  small  stream  into  the 
water,  not  the  water  into  the  acid. 

Sulphuric   acid  forms   several  hydrates  with  water, 

^the  two  most  important  being  those  represented  by  the 

formulas,  H2SO4.  H2O  and  H2SO4.  2  H3O.     Because  of 

the  tendency  of  sulphuric  acid  to  take  up  water,  it  is 

used  as  a  drying  agent. 

The  dehydrating  power  of  sulphuric  acid  also  accounts  for 
the  fact  that  organic  matter,  e.  </.,  paper,  wood,  dust,  sugar,  etc., 
are  charred  by  it.  These  bodies  are  compounds  containing 
carbon,  hydrogen,  and  oxygen  (cf.  §  52).  The  hydrogen  and 
oxygen  are  abstracted,  as  water,  by  the  acid,  and  charcoal 
is  left. 

187.  Reduction  of  Sulphuric  Acid.  —  Sulphuric  acid 
decomposes,   when   sufficiently  heated,  much  as  nitric 


178  SULPHUR  AND  ITS   COMPOUNDS. 

acid  does;  the  products  are  chiefly  sulphur  dioxide, 
oxygen,  and  water.  The  equation 

H2S04  =  H20  +  S02  +  O 

thus  represents  what  takes  place.  If  we  have  at  hand 
a  substance  that  can  absorb  oxygen,  the  decomposition 
of  the  acid  takes  place  very  easily.  The  reducing  agents 
may  be  sulphur,  charcoal,  copper,  mercury,  etc. 

The  acid  has  no  action  upon  these  substances  in  the  cold  ; 
but  when  it  is  heated  with  them  the  acid  is  reduced  to  sulphur- 
ous acid,  i.  e.,  to  sulphur  dioxide  and  water.  The  reducing 
agents  are,  of  course,  oxidized. 

The  following  equation  represents  what  takes  place 
when  sulphur  is  heated  with  sulphuric  acid  :  — 

2  H2SO4  4-  S  =  2  H2SO3  +  SO2  =  3  SO2  +  2  H2O. 
With  copper  the  equation  is,  — 

H2SO4  +  Cu  =  CuO  +  H2O  +  SO2. 

The  cupric  oxide  then  reacts  further  with  the  sulphuric 
acid,  giving  cupric  sulphate  and  water,  — 

CuO  +  H2S04  *=  CuS04  +  H20. 
Hence  the  complete  equation  is,— 

2  H2S04  +  Cu  =  CuS04  +  2  H2O  +  SO2. 

Perhaps  the  equations  for  the  action  of  copper  upon  hot, 
concentrated  sulphuric  acid  are,— 


SULPHATES.  179 

(1)  Cu  +  H2SO4  =  CuSO4  -f  H2,  and 

(2)  H2  -f  H2SO4  =  2  H2O  -f  SO2. 

In  any  case,  the  complete  equation  is  as  above, 

Cu  -f  2  H2SO4  =  CuSO4  +  2  H2O  -f  SO2. 

Compare  this  action  with  that  of  copper,  etc.,  upon  nitric 
acid  (§  159). 

1 88.  Uses  of   Sulphuric  Acid.  —  Sulphuric    acid  is 
used    in   many  processes    and    in    enormous   quantities. 
We  have  already  learned  that  it  is  used  in  the  preparation 
of  nitroglycerine   (cf.  §  165),  of  nitric  acid  (cf.  §  155), 
of  hydrochloric  acid  (cf.  §  91),  and  of  sodium  carbonate 
by  the  Le  Blanc  process  (cf.  §  92).     Sulphuric  acid  is 
used  also  in  the  refining  of  petroleum,  to  change  starch 
into  glucose,  and  to  convert  the  calcium  phosphate  of 
bone  ash  and  of  phosphate  rocks  into  soluble  form  for 
use  as  fertilizers. 

^N"o  wonder  that,  as  is  often  stated,  "  the  progress  of  civili- 
zation is  proportional  to  the  quantity  of  sulphuric  acid  used." 

189.  Sulphates.  —  The  salts  of  sulphuric  acid  may 
be    produced    in  the  usual  ways,   viz.,   by  neutralizing 
hydroxides  with  sulphuric  acid,  or  by  dissolving  metal- 
lic oxides,  metallic  carbonates,  or  the  metals  themselves 
in  the  acid. 

The  following  equations  represent  some  of  these  reactions  : 

2  NaOH  -f  H2SO4  =  Na2SO4  -f-  2  H2O. 
ZnCO3  -f  (dilute)  H2SO4  ==  ZnSO4  +  H2CO3. 
Mg  -f  (dilute)  H2SO4  =  MgSO4  -j-  H2. 


180  SULPHUR  AND  ITS   COMPOUNDS. 

Sulphates  may  also  be  formed  by  heating  chlorides, 
nitrates,  acetates,  etc.,  with  concentrated  sulphuric  acid 
(e/.  §§91  and  155). 

A  sulphate  that  is  difficultly  soluble  in  water  may  be 
formed  by  adding  dilute  sulphuric  acid  or  a  soluble  sul- 
phate to  a  solution  of  some  salt  of  the  metal.  Thus, 
strontium  sulphate,  SrSO4,  is  precipitated  Vhen  dilute 
sulphuric  acid  is  added  to  a  solution  of  strontium 
chloride. 

SrCl2  +  H2SO4  ==  /SVvS'04  +  2  HC1. 

Test. — -\Ve  usually  test  an  unknown  soluble  substance  to 
see  if  it  is  a  sulphate  by  adding  to  its  solution  barium  chloride 
solution  or  barium  nitrate  solution  ;  the  precipitate  of  barium 
sulphate  which  is  formed  if  a  sulphate  is  present,  is  insoluble 
in  acids. 

Thus,  BaCl2 -f  Xa28O4  =  #atf O4  +  2  NaCl. 

190.  Sulphur  Trioxide.  — Four  oxides  of  sulphur  are 
known;  but  only  the  trioxide  (SO3)  and  the  dioxide 
(8O2)  need  be  considered  here.  Sulphur  trioxide  is  a 
white,  crystalline  solid,  melting  at  15°  C.  and  very  solu- 
ble in  water.  When  it  is  brought  in  contact  with  water 
there  is  a  hissing  noise,  and  much  heat  is  liberated. 
The  product  of  the  union  is  sulphuric  acid. 

S03  +  H20  =  H2S04. 

Sulphur  trioxide  is  thus  the  anhydride  of  sulphuric 
acid. 

Sulphur  trioxide  is  produced  by  the  oxidation  of  sulphur  di- 
oxide ;  a  small  amount  is,  therefore,  formed  when  sulphur  is 


SULPHUR  DIOXIDE. 


181 


burned  in  air  or  in  oxygen.  It  is  formed  in  quantity  by  pass- 
ing a  mixture  of  dry  sulphur  dioxide  and  oxygen  over  heated 
platinum  sponge  or  platinized  asbestos. 

When  sulphur  trioxide  is  heated  considerably,  it 
breaks  up  in  part  into  sulphur  dioxide  and  oxygen. 

191.  Sulphur  Dioxide.  —  Sulphur  dioxide  is  the  gas 
of  Avell-known  odor  produced  by  burning  sulphur  in  air 
or  in  oxygen.  Commercially  it  is  usually  formed  by 
matting  sulphides,  e.  g.,  iron  pyrites,  in  air  (cf.  §§  178 
and  184).  The  reaction  is  represented  by  the  equation, 


240 


170 


=  Fe2O3 

IfiO 


4  SO 


From  its  power  of   striking   sparks  with  flint,    iron  pyrites 
derived  its  name  ;  u  pyrites  "  means  u  lirestone." 

The  method  commonly  em- 
ployed for  producing  sulphur 
dioxide  in  the  laboratory,  viz., 
by  the  reduction  of  concen- 
trated sulphuric  acid  with 
copper,  has  already  been  con- 
sidered (cf.  §  187). 

Sulphur  dioxide  is  a  color- 
less gas  about  2.2  times  as 
heavy  as  air.  One  liter  of  it 
Aveighs,  at  standard  condi- 
tions, 2.86  grams.  The  gas  is  very  soluble  in  water, 
80  c.c.  of  it  being  absorbed  by  1  c.c.  of  water  at  0°  C. 
When  the  solution  is  heated  the  gas  is  expelled. 


Ho  SO 4 

*    Mixture 

FIG.  42. 


K01I 


182  SULPHUK   AND   1TX   COMPOUNDS. 

Sulphur  dioxide  may  be  obtained  in  liquid  form  (Fig.  42)  by 
passing  it  through  a  condensing  tube  (conveniently  in  U  form) 
surrounded  by  a  freezing  mixture  of  ice  and  salt.  The  result- 
ing liquid  is  colorless  ;  it  boils  at  — 8°  C.  The  evaporation  of 
liquid  sulphur  dioxide  absorbs  much  heat ;  hence  this  sub- 
stance is  often  used  as  a  refrigerating  agent. 

192.  Chemical  Properties  of  Sulphur  Dioxide. — 

An  aqueous  solution  of  sulphur  dioxide  is  oxidized 
slowly  in  the  air  to  sulphuric  acid ;  the  oxidation  is  much 
more  rapid  if  oxidizing  agents  are  used. 

The  technical  preparation  of  sulphuric  acid  by  the 
oxidation  of  sulphur  dioxide  by  nitric  acid  has  been 
described  already  (<?/.  §  184).  Other  oxidizing  agents 
act  in  the  same  way. 

Thus,  solutions  of  potassium  chromate,  bichromate,  perman- 
ganate, etc.,  are  all  reduced  by  sulphur  dioxide.  The  latter  is 
converted  by  the  abstracted  oxygen  into  sulphuric  acid. 

The  chief  use  of  sulphur  dioxide,  aside  from  its  being 
the  source  of  sulphuric  acid,  is  as  a  bleaching  agent  of 
silks,  woolens,  straws,  laces,  etc.,  which  would  be  injured 
by  chlorine  (<?/.  §  88). 

The  sulphur  dioxide  probably  unites  with  the  coloring  matter 
of  these  fabrics,  forming  colorless  compounds.  The  bleaching 
effect  disappears  after  a  time,  however  ;  hence  fabrics  bleached 
by  sulphur  "  yellow  "  with  age.  Dilute  sulphuric  acid,  also, 
restores  the  color  of  many  articles  that  have  been  bleached  by 
sulphur  dioxide  ;  it  probably  breaks  up  the  colorless  compounds 
that  were  formed. 

Sulphur  dioxide  is  used  also  as  a  disinfectant. 


SULPHUROUS  ACID.  183 

193.  Sulphurous  Acid.  —  A  solution  of  sulphur  diox- 
ide in  water  has  the  properties  of  a  weak  acid,  and  is 
called  sulphurous  acid,  H2SO3.  Sulphur  dioxide  is, 
therefore,  sulphurous  anhydride.  The  free  acid  has  not 
been  isolated  (ef.  §'  149). 

When  sulphurous  acid  is  neutralized  by  the  solution 
of  a  base,  and  the  solution  is  evaporated,  a  sulphite  is 
obtained.  The  same  result  follows  when  sulphur  diox- 
ide gas  is  passed  into  the  solution  of  an  alkali. 

As  in  the  case  of  sulphuric  acid,  two  sets  of  salts  exist. 
Thus,  if  an  exactly  sufficient  amount  of  sulphur  dioxide  is  used 
with  sodium  hydroxide,  the  product  is  sodium  sulphite,  Na2SO3. 
The  equation  is,— 

HSO  -  »  NaSO        2  HO. 


2 


If,  however,  sulphur  dioxide  is  passed  into  sodium  hydroxide 
solution  until  no  more  gas  is  absorbed,  the  solution  contains 
sodium  hydrogen  sulphite,  NaHSO3.  The  equation  is,  — 


XaOII  +  H2SO8  -  »  NaIISO3  -f  H2O. 
Similar  reactions  take  place  with  other  bases. 

Both  normal  and  acid  sulphites  are  decomposed  by 
dilute  sulphuric  acid  and  by  hydrochloric  acid,  giving 
sulphur  dioxide.  Thus,  sodium  sulphite  and  hydro- 
chloric acid  react  according  to  the  equation, 

Na2SO8  +  2  HC1=  2  NaCl  +  H2SO8  (i.  <?.,  SO2  +  HaO). 

With  sodium  hydrogen  sulphite  the  equation  is, 
NaHSO8  +  HC1  =  NaCl  +  H2SO3  (*>?,  SO2  +  H2O), 


184  SULPHUR  AND  ITS   COMPOUNDS. 

By  this  reaction  a  sulphite  may  readily  be  distinguished 
from  a  sulphate. 

Sulphites,  like  sulphurous  acid,  oxidize  readily  in  the  air. 

194.  Thiosulphates.  —  Thiosulphuric  acid,  H2S2O3,  is 
sulphuric  acid  with  one-fourth  of  its  oxygen  replaced 
by  sulphur  ("  thion "  =  sulphur)  ;  its  salts  are  called 
thio sulphates. 

The  most  important  thiosulphate  is  the  sodium  salt, 
Na9S2O3.  This  is  made  by  boiling  a  solution  of  sodium 
sulphite  with  sulphur. 

Na2SO3  +  S  =  Na2S2O3. 

This  reaction  corresponds  to  the  oxidation  of  sulphites  to 
sulphates  (</.  §  193). 

When  sodium  thiosulphate  is  treated  with  dilute 
acids,  it  breaks  down  as  represented  by  the  equation, 

Na2S2O3  +  2  HC1  =  2  NaCl  +  H2SO3  +  S. 
It  decomposes,  therefore,  like  a  sulphite  plus  sulphur. 

Sodium  thiosulphate  is  a  reducing  agent  capable  of 
converting  chlorine  and  iodine  into  hydrochloric  and 
hydriodic  acids,  respectively.  It  is  therefore  used  to 
destroy  the  excess  of  chlorine  in  the  process  of  bleach- 
ing. It  has  the  power  to  dissolve  silver  chloride, 
bromide,  iodide,  etc.,  and  is  therefore  used  in  "fixing" 
negatives  in  photography.  Its  technical  name  is  "hypo," 
from  its  old  chemical  name,  "  hyposulphite  of  soda." 


EXERCISES.  185 

195.  Exercises. 

1.  Calculate   the  per  cent  of  sulphur  in  galena.     In  iron 
pyrites. 

2.  How  many  grams  of  sulphur  dioxide  can  be  produced  by 
burning  75  grams  of  sulphur?     How  many  liters  at  20°  C.  and 
740  mm.  ? 

3.  How  many  grams  of  ferrous  sulphide  are  required  to  give, 
with  an  excess  of  dilute  sulphuric   acid,  25  grams  hydrogen 
sulphide  ?     How  much  ferrous  sulphate  will  be  formed  ? 

4.  To  a  solution  containing  an  unknown  quantity  of  sul- 
phuric acid,  an  excess  of  barium  chloride  solution  was  added  ; 
the    resulting   barium    sulphate    weighed    2.653   grams.     How 
much  sulphuric  acid  was  there  in  the  solution  ? 

5.  At  least  how  many  grams  of  barium  chloride  crystals, 
BaCl2.  2  H2O,  are  required  to  produce  a  solution  that  will  pre- 
cipitate all   the  sulphuric  acid  formed  when  10  grams  of  sul- 
phur are  dissolved  in  fuming  nitric  acid? 

6.  How  could  you  distinguish  between  sodium  sulphite  and 
sodium  sulphate  ? 

7.  How  could  you  distinguish  between  sodium  chloride  and 
ammonium  chloride  ? 

8.  How  could  you  distinguish  between  a  soluble  sulphide, 
sulphite,  and  sulphate? 

9.  From  what  you  have  already  learned  of  ammonium  com- 
pounds, tell  what  will  probably  happen  when  you  heat  ammo- 
nium sulphide, 


CHAPTER   XIV. 
CARBON  AND   ITS   COMPOUNDS. 

196.  Carbon.  —  The  element  carbon  is  found  in  a 
free  and  an  almost  pure  state  as  diamond,  graphite,  and 
anthracite  coal  ;  it  is  found  combined  in  all  organic  sub- 
stances, in  carbon  dioxide,  in  carbonates,  in  coal,  and  in 
petroleum. 

Like  sulphur,  carbon  exists  in  several  allotropic  forms 
(cf.  §  175),  the  forms  usually  distinguished  being  (1) 
diamond,  (2)  graphite,  and  (3)  amorphous,  i.  e.,  non-crys- 
talline, carbon. 

These  three  modifications  of  carbon  differ  to  such  an  extent 
that  they  were  long  considered  different  chemical  individuals. 
Their  identity  is  proved  by  burning  them  in  oxygen  ;  for  equal 
parts  by  weight  of  each  give  equal  amounts  of  carbon  dioxide. 


12      32       44 

197.  The  Diamond.  —  The  diamond  is  prized  for  its 
luster,  its  strong  refractive  power,  and  its  hardness.  It 
is  one  of  the  hardest  substances  known. 

The  specific  gravity  of  the  diamond  is  3.5.  Acids 
and  alkalies  have  practically  no  effect  upon  it  ;  but 
when  it  is  heated  to  about  700°  C.  in  oxygen,  it  burns, 
forming  carbon  dioxide.  When  a  diamond  is  heated 


AMORPHOUS  CARBON.  187 

between  the  poles  of  a  powerful  battery  it  is  changed 
into  graphite. 

Diamonds  have  been  made  artificially  by  crystallizing  carbon 
from  solution  in  melted  iron.  This  probess  usually  gives  graph- 
ite, but  if  the  melted  iron  cools  under  great  pressure  the  carbon 
appears  in  the  form  of  small  diamonds.  The  necessary  pres- 
sure is  secured  by  chilling  the  exterior  of  the  mass  of  iron  ;  the 
contraction  of  the  exterior  thus  causes  great  pressure  upon  the 
interior  of  the  mass. 

So  far  as  known,  no  artificial  diamonds  yet  made  are  large 
enough  to  have  any  commercial  importance. 

198.  Graphite.  —  Graphite  is  sometimes  found  crys- 
tallized, but  usually  in  an  amorphous  form.     Its  spec- 
ific gravity  is  about  2.25.     It  owes  its  name,  "  black 
lead,"  to  a  confusion  of  names. 

Graphite  is  a  good  conductor  of  heat  and  of  electricity. 
Owing  to  its  friability  it  is  used  to  make  ulead  "  pencils. 
Mixed  with  clay  it  is  the  material  of  graphite  crucibles, 
which  are  used  in  making  "  crucible "  steel.  Other 
uses  are  :  To  protect  iron  from  the  air,  as  in  stove  polish, 
to  coat  grains  of  shot,  and  as  a  lubricant. 

Graphite  is  produced  artificially  (c/.  §  197)  by  crystallizing 
charcoal  from  molten  iron  and  steel. 

199.  Amorphous    Carbon.  —  The    amorphous   forms 
of  carbon  include  the  several  varieties  of  coal,  gas  car- 
bon, coke,   charcoal,  and  lampblack;   all    of  these  are 
formed  by  the  charring,  i.  e.,  carbonization,  of  animal  or 
vegetable  substances. 


188  CARBON  AND   ITS   COMPOUNDS. 

200.  Natural  Carbonization;  Coal. — Carbonization 
has  taken  place  in  nature  on  a  large  scale.     Vegetable 
matter,  accumulated  in  certain  localities  through  many 
generations,  and  decaying  in  the  absence   of    air,  was 
without  doubt  the  material  from  which  coal  was  formed. 
This  partially  decayed  matter  was  probably  much  like 
peat.     When  the  peat  was  buried  under  sediments,  and 
thus  subjected  to  water  and  pressure,  a  slow  carboniza- 
tion took  place;  as  a  result  gaseous  products,  such  as 
natural  gas,  etc.,  passed  off,  and  the  excess  of  carbon  re- 
mained behind. 

The  varieties  of  coal  owe  their  origin  to  the  different  degrees 
of  water-action  and  pressure  to  which  the  peat  was  subjected. 
Thus,  soft,  i.  e.,  bituminous,  coal  contains  many  gaseous  sub- 
stances, as  is  shown  by  its  burning  with  a  large  flame  and  much 
smoke  ;  anthracite  coal,  on  the  contrary,  is  nearly  pure  carbon, 
and  burns  with  a  small  flame. 

The  table  on  the  following  page  shoAVS  the  approximate 
composition  of  wood  and  of  several  varieties  of  coal. 
The  difference  between  the  sum  of  the  parts  per  cent 
and  100  represents  in  each  case  the  ash,  or  mineral  mat- 
ter, of  the  coal. 

201.  Artificial  Amorphous  Carbon.  —  The  artificial 
carbonization  of  wood  produces  wood  charcoal  and  lamp- 
black ;  that  of  soft  coal,  coke   and  gas  carbon;   that  of 
animal  matter  (bones,  blood,  etc.),  animal  charcoal  and 
lone-black. 


WOOD    CHARCOAL. 


189 


PER  CENT  OF 

CARBON. 

HYDROGEN. 

OXYGEN. 

NITROGEN. 

Wood. 

50 

6.0 

43 

Teat. 

62 

5.7 

32 

Lignite. 

68 

5.7 

25 

About  1. 

bituminous. 

79 

5.3 

14 

About  1. 

Cannel. 

83 

7.0 

9 

About  1. 

Anthracite. 

92 

3.5 

3 

About  1. 

Lampblack.  —  Lampblack,  or  soot,  is  produced  by 
burning  resinous  wood,  or,  for  the  better  grades,  oil  or 
gas,  in  an  insufficient  supply  of  air.  It  is  used  in  making 
printer's  ink,  black  paints,  India  ink,  etc. 

Wood  Charcoal.  —  Wood  charcoal  is  made  by  the 
''destructive  distillation  "  of  wood.  The  operation  may 
be  carried  out  either  in  iron  retorts,  or  by  piling  the 
wood  in  heaps  covered  with  sod  (Fig.  43)  and  then 
setting  fire  to  the  heaps.  In  the  latter  case  some  of  the 
wood  burns,  but  its  combustion  gives  heat  for  the  de- 
composition of  the  remainder. 

The  charcoal  obtained  from  a  given  mass  of  wood  weighs 
from  15%  to  25%  as  much  as  the  wood  taken  ;  the  loss  is 
due  to  the  escape  of  volatile  substances.  The  gaseous  products 


190 


CAEBON  AND  ITS   COMPOUNDS. 


were  formerly  used  as  illuminating  gas  ;  the  liquid  parts  con- 
sist of  wood  spirit  (methyl  alcohol),  acetic  acid,  etc. 


Wood  charcoal  is  used  in  the  reduction  of  iron  from  its 
ores  and  as  a  disinfectant.  It  absorbs  large  quantities 
of  certain  gases :  ninety  volumes  of  ammonia,  or  nine 
volumes  of  oxygen  are  known  to  have  been  taken  up 
by  one  volume  of  box-wood  charcoal.  Charcoal  owes 
its  action  as  a  disinfectant  to  its  power  of  absorb- 
ing noxious  gases,  along  with  oxygen,  in  its  pores. 
The  oxygen  destroys  the  bacteria  in  the  other  gases. 

Coke  and  Gas-Carbon.  —  Coke  is  the  residue  left 
when  soft  coal  is  distilled;  gas-carbon  is  the  volatile 
coke  condensed  upon  the  walls  of  the  retorts. 

Gas-carbon  is  metallic  in  character  and  a  good  con- 
ductor of  electricity;  it  is  used  for  the  negative  plates 
of  electric  cells  and  for  the  pencils  of  electric  arc-lamps. 

Coke  is  used  chiefly  in  metallurgy  to  reduce  the 
metals  from  their  ores.  It  is  also  a  fuel. 


CARBON  DIOXIDE.  191 

The  volatile  products  obtained  in  the  distillation  of  coal  con- 
sist of  illuminating  gas,  coal-tar  (the  source  of  many  organic 
compounds),  ammonia,  etc.  (c/.  §  143). 

Animal  Charcoal  and  Bone-black.  —  Animal  char- 
coal and  bone-black  are  used  to  destroy  organic  coloring 
matter.  When  a  solution  of  brown  sugar,  for  example,  is 
filtered  through  bone-black,  the  solution  becomes  color- 
less. Vinegar  may  be  clarified  in  the  same  way. 

When  bones  are  distilled  destructively  the  volatile  products 
consist  largely  of  bone-oil.  Usually  the  bones  are  deprived  of 
their  gelatine  before  being  carbonized.  In  either  case  the  ani- 
mal charcoal  obtained  contains  the  mineral  matter  of  the 
bones. 

202.  Carbon  Dioxide  ;  Occurrence. — There  are  two 
compounds  of  carbon  and  oxygen,  viz.,  carbon  monoxide 
and  carbon  dioxide  ;  of  these  the  latter  is  by  far  the  more 
important. 

Carbon  dioxide  is  present  in  the  air,  in  some  mineral 
springs,  and  in  certain  localities  where  it  escapes  from 
the  earth.  Ten  thousand  parts  by  volume  of  air  con- 
tain, on  the  average,  about  3.5  parts  of  carbon  dioxide 
(ff.  §  117). 

Carbon  Dioxide  Exhalations.  —  At  Herste,  near  Driburg,  Ger- 
many, certain  borings  made  in  1894  struck  carbon  dioxide  at  a 
depth  of  148.5  meters.  A  violent  outburst  of  the  gas  took 
place,  no  less  than  40,000,000  liters  escaping  daily.  The  car- 
bon dioxide  was  99.84%  pure.  In  1897,  10,000  kilograms,  i.e., 
about  one-eighth  of  the  outflow,  were  liquefied  daily. 

The  origin  of  the  carbon  dioxide  was  probably  the  action  of 


192  CARBON  AND  ITti   COMPOUNDS. 

certain  silicates,  i.  e.,  salts  of  silicic  acid,  H2SiO3,  upon  calcium 
carbonate. 

203.  Preparation  of  Carbon  Dioxide.  —  In  the  lab- 
oratory, carbon  dioxide  is  commonly  made  by  decompos- 
ing a  carbonate  by  an  acid.     The  carbonate  generally 
used  is  marble,  CaCO3.     The  reaction  of  marble  with 
hydrochloric  acid  is  represented  by  the  equation, 

CaC03  +  2  HC1  =  CaCl2  +  H2CO8  (i.  e.,  CO2  +  H2O). 

Evaporation  of  the  solution,  after  all  action  ceases,  gives 
calcium  chloride,  CaCl2,  the  substance  used  to  dry  gases,  etc. 

204.  Physical    Properties.  —  Carbon    dioxide    is    a 
colorless  gas.     It  has  the  slightly  acid  taste  known  to  all 
who  drink  "  soda  water."*  Jt*is  about  one  and  one-half 
times  as  heavy  as  air ;  one  liter  weighs  1.97  grams  at 
standard  conditions. 

Gaseous  carbon  dioxide  may  be  condensed  to  the 
liquid  state  at  ordinary  temperatures  by  -a  pressure  of 
about  fifty  atmospheres  ;  above  31°  C.,  however,  —  its 
critical  temperature,  —  the  gas  cannot  be  liquefied  by 
pressure.  If  liquid  carbon  dioxide  is  allowed  to  evapo- 
rate rapidly,  it  solidifies,  forming  a  white  mass  like 
snow.  When  liquid  air  (cf.  §  122)  evaporates,  solid 
carbon  dioxide  is  usually  left  behind. 

About  ten  million  kilograms  of  liquid  carbon  dioxide  are 
used  annually. 

Water  absorbs  about  1.8  times  its  OAVH  volume  of 
carbon  dioxide  gas  at  standard  conditions  ;  with  increase 


CHEMICAL   PROPERTIES.  193 

of  pressure  the  amount  dissolved  increases.  When  the 
excess  of  pressure  is  removed  from  water  surcharged 
with  carbon  dioxide,  much  carbon  dioxide  escapes ; 
hence  the  sparkling  of  "  soda  water  "  and  also  of  spring 
water,  which  usually  has  carbon  dioxide  in  solution. 

205.  Chemical  Properties.  —  Carbon  dioxide  does 
not  allow  ordinary  burning  to  continue  in  it  any  more 
than  water  does,  and  for  the  same  reason. 

Because  of  its  inability  to  support  combustion,  carbon  di- 
oxide is  used  to  extinguish  fires.  For  this  purpose  it  is  gen- 
erated in  a  strong  reservoir  by  the  action  of  dilute  sulphuric 
acid  upon  sodium  carbonate. 


:Na2CO3  -f  H2SO4  =  Na2SO4  +  H2CO3  (i.  e.,  H2O  -f  CO2). 
The  resulting  gas  is  directed  in  a  stream  upon  the  flames. 


FIG.  44. 


Very  strongly  burning  substances,  however,  continue 
to  burn  in  the  gas;  examples  are  sodium  and  mag- 
nesium. 


194  CARBON  AND  IT  8   COMPOUNDS. 

The  equations  for  the  action  of  burning  sodium  upon  carbon 
dioxide  are,  — 

(1)  4  Na  +  C02  =  2  Na20  +  C. 

(2)  Na2O  +  CO2  =  Na2CO3. 


When   carbon  dioxide  is  passed  over  red-hot  carbon 
(Fig.  44)  it  is  reduced  to  carbon  monoxide  (cf.  §  212). 


C02  +  C  -  .200. 


The  volume  of  the  carbon  dioxide  formed  by  the 
union  of  a  given  volume  of  oxygen  with  carbon  is 
equal  to  the  volume  of  the  oxygen. 


12       32  44 

1  vol.     1  vol. 

Of  the  volume  of  the  carbon  nothing  can  be  stated,  because 
carbon  cannot  readily  be  obtained  in  the  gaseous  state. 

206.     Other  Sources  and  Uses  of  Carbon  Dioxide.  - 
.  _  (a)  Fermentation.    When 

an  aqueous  solution  of  cane 
sugar  or  grape  sugar  is  sub- 
jected to  the  action  of  the 
yeast  plant,  the  sugar  is  de- 
composed ;  the  chief  products 
are  ethyl  alcohol,  C2H5OH, 

and    carbon    dioxide.       The 
FIG.  45. 

alcohol  remains    in   solution, 

but  most  of  the  carbon  dioxide  escapes  in  gaseous  form. 


BAKING   POWDER 8.  195 

Fig.  45  shows  some  characteristic  yeast  plants  in  the  pro- 
cess of  budding. 

The  decomposition  of  sugar  by  yeast  is  called  alco- 
holic fermentation.  Grape  sugar  is  a  complex  compound 
of  the  formula  C6H12O6 ;  its  decomposition  in  fermen- 
tation is  chiefly  according  to  the  equation, 

C6H1306  =  2C03+2C3H6OH. 

The  brewer  uses  fermentation  to  produce  alcoholic 
liquors ;  the  baker,  to  raise  bread. 

(6)  Baking  Powders.  Ordinary  baking  powder  is 
a  mixture  of  potassium  hydrogen  tartrate,  KHC4H4O6 
(commonly  called  "  cream  of  tartar  "),  and  sodium  bi- 
carbonate, NaHC()3,  with  some  diluting  substance,  e.  g. 
corn-starch.  The  ingredients  act  upon  each  other  only 
when  moist  or  in  solution.  The  equation  for  the  de- 
composition of  baking  powder  is,  — 

KHC4H4O6  +  NaHCO3  =  KNaC4H4O6  +  H2CO3. 
The  substance  KNaC4H4O6  is  called  "  Rochelle  Salt." 

Theoretically,  any  substance  that  will  liberate  carbon  diox- 
ide from  sodium  bicarbonate  might  be  used  in  place  of  cream 
of  tartar,  but  practically  a  substance  must  be  used  that  will 
not  form  a  residue  injurious  to  the  human  system. 

"  Acid  Phosphate "  baking  powders  contain  so- 
dium bicarbonate  and  calcium  hydrogen  phosphate, 
CaH4(PO4)2.  The  calcium  hydrogen  phosphate  liber- 


196        CARBON  AND  ITS  COMPOUNDS. 

ates    carbon  dioxide  from  sodium  bicarbonate,   just  as 
cream  of  tartar  does. 

207.  Relation  of  Carbon  Dioxide  to  Life.  —  Sub- 
stances containing  carbon  usually  give  by  their  oxida- 
tion carbon  dioxide.  The  same  result  follows  whether 
the  oxidation  is  slow  or  rapid.  Hence  all  decay,  as  of 
wood,  paper,  etc.,  results  in  the  formation  of  this  gas. 

Carbon  dioxide  fails  to  support  animal  life,  not  because 
carbon  dioxide  is  poisonous,  but  because  animals  cannot  ex- 
tract from  it  the  oxygen  necessary  for  respiration.  Besides. 
the  presence  of  even  a  little  more  than  the  normal  amount  of 
carbon  dioxide  in  the  air,  say,  7  or  8  parts  in  10,000,  prevents 
the  proper  exit  of  carbon  dioxide  from  the  lungs. 

The  fact  that  the  enormous  quantity  of  carbon 
dioxide  constantly  poured  into  the  air  by  the  respiration 
of  animals  and  plants  does  not  accumulate  until  it 
destroys  higher  animal .  life  is  due  to  the  agency  of 
vegetation. 

To  plants,  carbon  dioxide  is  an  important  food ;  for 
chlorophyll-producing,  i.  e.,  green,  plants  are  able  to  con- 
vert carbon  dioxide  and  water,  m  the  presence  of  light, 
into  sugar,  starch,  wood,  and  the  various  compounds  of 
carbon,  hydrogen,  and  oxygen  of  which  most  vegetable 
products  consist. 

The  natural  cycle  through  which  carbon  dioxide  passes  is 
impressive  :  — 

(1)  Plants,  by  means  of  energy  derived  from  the  sun, 
change  carbon  dioxide  and  water  into  plant  tissue; 


CARBONATES.  197 

(2)  Animals,  appropriating  the  stored-up  energy  of  plants, 
give  it  forth  in  their  activities  and  return  carbon  dioxide  to 
the  air. 

208.  Carbonic  Acid.  —  The  aqueous  solution  of 
carbon  dioxide  has  slightly  acid  properties  and  gives 
carbonates  with  soluble  bases  ;  hence  the  substance 
H9CO3  —  carbonic  acid  —  probably  exists  in  solution. 
Carbon  dioxide  is  thus  carbonic  anhydride. 

Carbonic  acid  cannot  be  isolated  ;  for  it  breaks  up 
very  readily  into  carbon  dioxide  and  water. 


In  its  instability  carbonic  acid  is  like  ammonium  hydroxide 
and  sulphurous  acid  (cf.  §§  149  and  193). 

209.  Carbonates.  —  When  carbon  dioxide  is  passed 
into  solutions  of  metallic  hydroxides  it  is  absorbed, 
forming  carbonates. 

Thus,  sodium  hydroxide  and  carbonic  acid  (carbon  dioxide 
with  water  is  probably  carbonic  acid)  react  according  to  the 
equation, 


2  NaOTI  +  H2C03  (i.  e.  ,  CO,  +  IT/))  =  Xa/1O3  +  2  H2O. 

A  similar  reaction  takes  place  when  carbon  dioxide  is 
passed  into  a  solution  of  calcium  hydroxide  (lime- 
water)  ;  but  in  this  case  the  carbonate  formed  is  in- 
soluble. Asa  result  the  lime-water  becomes  "  milky." 

The  equation  is,  — 

Ca(OH)2  +  H2CO3  :=;  Ca  CO^  +  2  H2O. 


198        CARBON  AND  ITS  COMPOUNDS. 

With  baryta-water  (barium  hydroxide  solution)  the  result  is 
analogous. 

Ba(OH)2  +  H2C03  =  BnCOa  +  2  H2O. 

If  a  gas  forms  a  white  precipitate  on  being  passed  into 
lime-water  or  baryta-water,  we  generally  assume  that  the 
gas  is  carbon  dioxide  (cf.  §  117). 

All  the  common  carbonates  except  those  of  sodium, 
potassium,  and  ammonium  are  insoluble  in  pure  water, 
and  are,  therefore,  precipitated  when  one  of  the  three 
soluble  carbonates  is  added  to  the  solution  of  a  metal 
salt. 

Thus,  sodium  carbonate  gives  with  calcium  chloride  a  pre- 
cipitate of  calcium  carbonate. 

CaCl2  -f  Xa2CO3  =  CaC03  +  2  XaCl. 

Most  carbonates  except  those  of  sodium  and  potas- 
sium decompose,  when  heated,  into  the  corresponding 
oxides  and  carbon  dioxide. 

Thus,  marble  gives,  when  heated,  quicklime  and 
carbon  dioxide. 

CaCO 


Ammonium  carbonate  is,  of  course,  decomposed  by  heat,  like 
other  ammonium  salts.  It  is  often  called  iks«Z  volatile." 

The  common  method  of  identifying  a  carbonate  is  to 
treat  it  with  a  dilute  acid  and  then  to  prove  that  the 
escaping  gas  is  carbon  dioxide. 

210.  Bicarbonates.  —  If  carbon  dioxide  is  passed  into 
lime-water  long  enough,  the  precipitate  of  calcium  car- 


BICAEBONATES.  199 

bonate,  which  is  formed  at  first,  will  be  redissolved.  A 
similar  result  takes  place  with  barium  hydroxide.  This 
phenomenon  may  be  best  understood  by  comparing  it 
with  the  action  of  carbonic  acid  (i.  e;,  carbon  dioxide 
and  water)  upon  sodium  carbonate  —  a  reaction  repre- 
sented by  the  equation, 

Na2C03  +  H3C08  (i.  e.,  CO2  +  H2O)  -  »  2  NaHCO3. 


The  substance  NaHCO3  is  sodium  hydrogen  carbonate  (called, 
also,  sodium  bicarbonate)  ;  when  it  is  heated  gently  it  breaks 
up  into  sodium  carbonate,  carbon  dioxide,  and  water. 

2  NaHCOg  -  »  NajjCOg  +  HaCO,  (L  e.,  H3O  +  CO,). 


In  a  similar  way  an  excess  of  carbonic  acid  converts 
calcium  carbonate  into  calcium  hydrogen  carbonate, 
Ca(HCO3)2.  The  precipitate  of  calcium  carbonate  re- 
dissolves,  therefore,  owing  to  the  formation  of  calcium 
bicarbonate,  which  is  soluble.  The  equation  is,  — 

CaC03  +  H2C08  -  »  Ca(HC03)2. 

Calcium  bicarbonate  is,  however,  very  unstable  ;  it  de- 
composes, when  its  solution  is  heated,  into  calcium 
carbonate,  carbon  dioxide,  and  water.  Hence  the  precipi- 
tate of  calcium  carbonate  reappears  on  boiling. 

The  behavior  of  calcium  carbonate  with  an  excess  of  car- 
bon dioxide  explains  the  solubility  of  limestone  in  natural 
waters  which  contain  this  gas  ;  and  the  ease  with  which  cal- 
cium bicarbonate  is  decomposed  explains  why  most  waters 
deposit  their  limestone  upon  the  walls  of  vessels  in  which  they 
are  heated  (c/.  §  40). 


200  CARBON  AND  ITU   COMPOUNDS. 

In  many  limestone  regions  underground  waters  are 
charged  under  pressure  with  carbon  dioxide,  and  there- 
fore take  up  large  quantities  of  limestone  ;  subsequently 
the  carbon  dioxide  escapes,  and  the  limestone  is  de- 
posited. These  limestone  deposits  often  take  on  beau- 
tiful and  fantastic  forms,  as  stalactites,  etc.  See  Fig.  46. 


FIG.  46. 
CAVE  AT  MARENGO,  INDIANA. 

211.    Natural    Carbonates.  —  The    most    abundant 
natural   carbonate   is    limestone,   CaCOg.     Large    quan- 


FORMATION  OF   CARBON  MONOXIDE.  201 

titles  of  this  substance  are  distributed  over  the  land  and 
in  the  sea. 

Marble  is  a  finely  crystallized  limestone.  Iceland  spar  is 
almost  pure  calcium  carbonate  ;  it  occurs  in  large  crystals. 
Coral  is  limestone  taken  from  the  sea  by  the  coral  polyp  and 
left  as  a  residue  when  the  polyp  dies.  The  shells  of  most  water 
animals  are  largely  limestone. 

A  natural  mixture  of  calcium  and  magnesium  carbon- 
ates called  dolomite  is  also  very  abundant ;  it  is  the  chief 
component  of  whole  ranges  of  the  Alps  mountains. 
Magnesite  is  natural  magnesium  carbonate. 

Sodium  carbonate  and  potassium  carbonate  exist  in  the 
soil  of  many  places.  The  ashes  of  sea-plants  contain 
sodium  carbonate ;  those  of  land-plants,  potassium  car- 
bonate. 

Both  of  these  carbonates  are  used  in  large  quantities  and 
are  made  synthetically,  especialry  sodium  carbonate,  on  an 
enormous  scale  (cf.  §  92).  A  large  and  valuable  deposit  of 
almost  pure  sodium  carbonate  occurs  at  Owen's  Lake,  Cali- 
fornia. 

212.  Formation  of  Carbon  Monoxide.  —  Carbon 
monoxide  is  produced  by  the  reduction  of  carbon 
dioxide  by  hot  carbon.  This  operation  is  illustrated  in 
Fig.  44.  The  equation  is,  — 

CO2  +  C »2CO. 

A  common  illustration  of  the  same  reaction  is  seen  in  every 
coal  fire.  The  coal  at  the  bottom  (Fig.  47)  burns  to  carbon 
dioxide  ;  but  as  this  gas  passes  through  the  bed  of  heated  coal 


202 


CAEBON  AND  ITS   COMPOUNDS. 


it  is  reduced  to  carbon  monoxide.     The  carbon  monoxide  then 
burns  at  the  top  of  the  bed  of  coal  to  carbon  dioxide. 


2  =  2  CO2. 


213.  Laboratory  Method.  — 

The  common  method  of  pre- 
paring carbon  monoxide  is  to 
heat  crystallized  oxalic  acid, 
H2C2O4.  2  H2O,  with  concen- 
trated sulphuric  acid. 

The  purpose  of  the  sulphuric 
acid  is  to  take  up  the  water  of 
the  oxalic  acid,  not  only  the 
water  of  crystallization,  but  also 

the  water  which  the  oxalic  acid  gives  when  it  decom- 

poses. 


FIG.  47. 


H2C204  =  H20  +  CO  +  C0 


2. 


The  carbon  monoxide  is  freed  from  carbon  dioxide  by 
passing  the  mixed  gases  through  sodium  hydroxide  solu- 
tion. The  sodium  hydroxide  combines  with  the  carbon 
dioxide  (cf.  §  209),  but  not  with  the  carbon  monox- 
ide. 

Another  method  of  making  carbon  monoxide  is  to  heat 
potassium  ferrocyanide,  K4Fe(CN)6  (called,  also,  "  yellow  prus- 
siate  of  potash"),  with  concentrated  sulphuric  acid. 


214.  Properties.  —  Carbon  monoxide  bears  to  formic 
acid,  H2CO2,  a  relation  similar  to    that  which   carbon 


CYANOGEN.  203 

dioxide  bears  to  carbonic  acid.     This  is   evident  from 
the  formulas. 

Formic  acid,  H2CO2.     Carbon  monoxide,  CO. 

Carbonic  acid,  H2CO3.     Carbon  dioxide,  CO2. 

Water  and   carbon    monoxide    do    not  unite,   however, 
under  ordinary  conditions. 

Carbon  monoxide  is  a  colorless  gas  ;  when  pure  it  is 
almost  odorless.  It  is  popularly  known  as  "  coal  gas" 
and  is  familiar  to  every  one  who  has  used  anthracite 
coal.  It  burns  with  a  beautiful,  blue  flame.  The  flame 
is  best  observed  when  fresh  coal  is  put  upon  a  hot  fire. 

The  flame  of  carbon  monoxide,  like  that  of  hydrogen,  is 
simple,  i.  e.,  it  consists  of  only  one  region  of  combustion  (cf. 
§11). 

The  equation  for  the  combustion  of  carbon  monoxide 
is,  — 


2  vols.     1  vol.     2  vols. 

Unlike  carbon  dioxide,  carbon  monoxide  is  extremely 

poisonous  ;  air  containing  one  per  cent  of  it  will  pro- 
duce fatal  results. 

One  liter  of  carbon  monoxide  weighs  1.25  grams  at 

standard  conditions.  The  gas  has  been  liquefied  at 

—141°  C.,  its  critical  temperature,  by  a  pressure  of 
35  atmospheres. 

215.  Cyanogen.  —  Carbon  and  nitrogen  combine  to 
form  a  gas  called  cyanogen,  C2N2.     Cyanogen  may  be 


204  CARBON  AND   ITS   COMPOUNDS. 

prepared  by  heating  mercuric  cyanide,  Hg(CN)2,  or  sil- 
ver cyanide,  AgCN. 

With  silver  cyanide  the  equation  is,  — 


A  cheaper  way  is  to  allow  a  solution  of  potassium  cyanide, 
KdN",  to  fall  drop  by  drop  into  ahot  solution  of  eupric  sulphate. 
Cupric  cyanide,  Cu(CN")2,  is  first  formed,  but  at  once  breaks  up 
into  cuprous  cyanide,  Cu2(dN")2,  and  cyanogen. 

(1)  CuS04  +  2  KCN  =  Cu(CN)2  +  K2SO4. 

(2)  2 


Cyanogen  is  a  colorless,  sweet-smelling  gas.  It  burns 
with  a  beautiful,  lavender-colored  flame,  forming  carbon 
dioxide  and  nitrogen. 


The  flame  of  cyanogen  is  double,  i.  e.,  it  has  two  zones  of  com- 
bustion (c/.  §§  11  and  214). 

Cyanogen  is  so  poisonous  that  no  one  but  an  experi- 
enced person  should  make  it  in  quantity,  and  then  only 
where  there  is  a  strong  draught. 

216.  Hydrocyanic  Acid.  —  Hydrocyanic,  or  prussic, 
acid  is  extremely  poisonous.  It  may  be  made  by  treat- 
ing cyanides,  e.  g.,  potassium  cyanide  or  potassium  ferro- 
cyanide,  with  dilute  sulphuric  acid. 

KCN  +  H2SO4=  KHSO4  +  HCN. 
It  is  a  colorless,  sweet-smelling  liquid,  boiling  at  27°  C, 


METHANE.  205 

217.  Compounds  of  Carbon  and  Hydrogen.  —  The 

number  of  compounds  formed  by  carbon  and  hydrogen 
is  very  large.  At  present  we  shall  consider  only  four, 
viz.  :  methane,  CH4 ;  ethane,  C2H6 ;  ethylene,  C2H4,  and 
acetylene,  C9H9.  All  are  colorless  gases. 

218.  Methane.  —  Methane,  or  marsh  gas,  is  formed 
in  the  destructive  distillation  of  coal,  and  is  therefore 
present  in  illuminating  gas  made  by  the  old  process  (/:/'. 
§  223).     It  is  formed  in  nature  by  the  decay  of  organic 
matter,  e.  //.,  leaves,  twigs,  etc.,  under  water ;  hence  the 
name  "marsh  gas."     The  presence  of  the  gas  may  be 
observed   by  disturbing  the  material  at  the  bottom  of 
most  stagnant  pools,  especially  in  summer. 

Marsh  gas  also  enters  coal  mines  ;  here  its  mixture  with  air 
forms  the  dreaded  "fire  da-nip  "  ( damp  =  gas  ;  cf.  Ger.  Dampf.), 
which  explodes  with  frightful  violence  when  ignited.  It  was 
to  avoid  these  explosions  in  mines  that  Davy  made  his  cele- 
brated "  safety  lamp  "  (</.  §  32). 

Methane  forms  over  ninety  per  cent  of  the  gases  that 
escape  from  petroleum  wells. 

Marsh  gas  is  commonly  made  in  the  laboratory  by 
heating  a  mixture  of  sodium  acetate,  NaC9H3O9,  sodium 
hydroxide,  and  quicklime.  The  gas  thus  produced  is 
not  pure,  however ;  it  contains  hydrogen  and  other  im- 
purities. An  approximate  equation  is, 

NaC2H3O2  +  NaOH  =  Na2CO3  +  CH4. 


206  CARBON  AND  ITS   COMPOUNDS. 

The  equation  representing  the  combustion  of  marsh 
gas  is,  — 


1  vol.      2  vols.       1  vol.      2  vols.  (steam) 

219.  Ethane.  —  Ethane  is  formed  along  with  marsh 
gas  in  the  decomposition  of  many  organic  substances, 
and  is  a  constituent  of  illuminating  gas.     Its  composi- 
tion is  represented  by  the  formula,  C2H6.     It  burns  with 
an  almost  colorless  flame. 

In  burning,  ethane  produces  carbon  dioxide  and  water,  just 
as  marsh  gas  does,  but  the  relative  proportions  are  different. 
This  fact  is  shown  by  the  equation, 

2  C2H6  -f  7  O2  =  4  CO,  -f  G  H2O. 

2  vols.  7  vols.     4  vols.       6  vols.  (steam) 

220.  Ethylene.  —  Ethylene  may  be  prepared  by  heat- 
ing ethyl  alcohol,  C2H5OH,  with  concentrated  sulphuric 
acid  to  above  140°  C.     The  following  equation  repre- 
sents in  part  what  takes  place  :  — 

C2H5OH  —  H2O  ==  C2H4. 

Ethylene  is  present  in  illuminating  gas  ;  it  burns  with 
a  bright,  luminous  flame,  producing  carbon  dioxide  and 
water. 

C2H4  +  3  02  =  2  C02  +  2  H20. 

1  vol.         3  vols.      2  vols.         2  vols. 

The  most  remarkable  property  of  ethylene  is  its  power  to 
absorb  chlorine  and  bromine,  giving  ethylene  chloride  and 
bromide. 


ACETYLENE.  207 

Thus,  C2H4  +  2  Br  =  C2H4Br2. 

28         160  188 

Both  the  chloride  and  the  bromide  are  heavy,  colorless  oils. 

221.  Acetylene.  —  Like  the  other  compounds  of 
hydrogen  and  carbon  that  have  been  mentioned,  acety- 
lene is  present  in  illuminating  gas.  It  is  formed  when 
a  Bunsen  burner  burns  at  the  base  instead  of  at  the  top, 
and  may  then  be  recognized  by  its  penetrating  odor. 

Acetylene  has  been  formed  by  passing  an  electric 
spark  between  carbon  terminals  in  a  vessel  of  hydrogen. 

2  C  +  H2  =  C2H2. 

Acetylene  has  recently  been  made  on  a  commercial  scale  by 
the  action  of  water  containing  a  little  hydrochloric  acid  upon 
the  carbides  of  certain  metals.  The  carbide  commonly  used  is 
calcium  carbide,  CaC2.  The  decomposition  of  this  substance 
by  water  is  represented  by  the  equation, 

CaC2  +  2  H20  =  C2H2  +  Ca(OH)2. 

Calcium  carbide  is  produced  by  heating  a  mixture  of  quick- 
time,  CaO,  and  coal,  or  coke,  in  the  electric  furnace.  The 
reaction  is  represented  thus  :  — 

CaO  +  30  =  CaC2  +  CO. 
A   section    of   a   simple 
electric  furnace  is 
in  Fig.  48. 


Acetylene  burns, 
under    proper   condi- 
tions, with  a  flame  of  FIG-  48- 
dazzling  whiteness  and  without  soot.     It  is  readily  con- 


208       CAEBON  AND  ITS  COMPOUNDS. 

densed  to  the  liquid  state.    Ordinary  mixtures  of  gaseous 
acetylene  and  air  are  very  explosive. 

The  equation  for  the  combustion  of  acetylene  is,  — 

2  C2H2  +  5  02  ==  4C02  +  2  II2<). 
2  vols.         5  vols.       4  vols.          2  vols. 

222.  Illuminating  Gas.  — Illuminating  gas  is  a  mix- 
ture of  several  gases  already  studied.     These  gases  may 
be  divided  into  (#)  combustible  gases  and  (7>)  impurities. 

The  impurities  are  chiefly  nitrogen  find. carbon  dioxide. 

The  combustible  gases  are  of  two  kinds  :  - 

(1)  Those  that  burn  without  giving  light,  and  (2) 
those  that  burn  with  luminous  flames. 

The  non-illuminating  gases  are  hydrogen,  methane,  and 
carbon  monoxide ;  they  make  up  about  ninety  per  cent 
by  volume  of  the  gas.  The  illuminating  gases  are  the 
so-called  "  heavy  hydrocarbons,"  ethane,  propane,  03H8, 
butane,  C4H10,  etc.,  and,  usually,  small  amounts  of 
ethylene  and  acetylene. 

Two  general  processes  are  employed  in  making  illumi- 
nating gas  :  — 

(1)  The  distillation  of  soft  coal. 

(2)  The  "  water-gas  "  process. 

223.  Illuminating  Gas  by  Distillation  of   Coal. - 

The  old  process  of  making  gas  is  carried  out  as  shown 
in  Fig.  49. 

Soft  coal  in  the  retorts  C  (there  are  usually  several  retorts, 
one  over  the  other)  is  heated  by  the  fire  A  to  the  temperature 


ILL  UMINA  TING  GAS  B  Y  DIti  TILL  A  TION  OF  COAL.     209 

of  decomposition.  The  volatile  products  pass  off  through  the 
pipe  T  into  the  "  hydraulic  main  "  B.  The  hydraulic  main 
contains  water,  which  condenses  much  tar,  etc.,  from  the  gas. 
From  B  the  gas  passes  through  the  "  condensers  "  D,  which 
stand  over  water  ;  here  the  gas  is  cooled,  and  more  of  the 
tarry  products  are  condensed. 


FIG.  49. 

From  the  condensers  the  gas  passes  through  the  coke  towers, 
or  "  scrubbers  "  O  ;  into  these  water  is  sprayed,  and  the 
illuminating  gas  is  thereby  freed  from  soluble  gases,  such  as 
ammonia  and  hydrogen  sulphide. 

From  the  "  scrubbers  "  the  gas  passes  into  the  "  purifiers  " 
M.  The  purifiers  are  large  boxes  containing  trays  of  slaked 
lime  ;  this  substance  removes  carbon  dioxide  and  traces  of 
hydrogen  sulphide.  The  gas  then  enters  the  gas  holder  G  ; 
thence  it  is  distributed  to  the  community  through  the  service 
pipe  S'. 


Many  other  valuable    products    besides   illuminating 


210  CARBON  AND   ITS   COMPOUNDS. 

gas  are  obtained  by  the  distillation  of  coal.  The  am- 
moniacal  liquors  of  the  condensers  arid  coke  towers  are 
the  sources  of  ammonium  compounds  (</.  §  143)  ;  the 
tar  of  the  hydraulic  main  and  condensers  gives,  when 
distilled,  the  important  substances  benzene,  toluene, 
phenol  (carbolic  acid),  naphthalene,  etc. ;  the  residue  in 
the  retorts  is  coke  and  gas  carbon  (cf.  §  201). 

The  composition  of  a  representative  illuminating  gas  made 
by  the  old  process  is  shown  in  the  table  accompanying  §  225. 

224.  Water-gas  Process.  —  Within  recent  years  the 
old  process  for  the  production  of  artificial  gas  has  been 
almost  entirely  displaced  by  the  "  water-gas  "  process. 
In  this  process  steam  is  passed  over  white-hot  anthracite 
coaL  The  reactions  which  take  place  are  represented 
*by  the  equations, 

(1)  C  +  2  H20  =*=  C02  +  2  H2. 

(2)  CO2  +  C  =  2  CO." 

The  mixture  of  carbon  monoxide  and  hydrogen  is 
"  water-gas."  It  may  be  used  directly  for  fuel,  but  not 
for  lighting,  since  it  burns  with  a  colorless  flame. 

For  the  conversion  of  water-gas  into  an  illuminating 
gas,  the  water-gas  is  "  enriched  "  by  the  addition  of 
petroleum  hydrocarbons.  The  process  of  making  an 
illuminating  water-gas  will  be  understood  from  Fig.  50. 

Anthracite  coal  is  heated  to  white  heat  in  the  "generator  " 
by  means  of  a  blast  of  air  forced  through  the  coal  by  the 
engine.  The  blast  of  air  is  then  cut  off,  and  superheated 
steam  is  forced  ~»ver  the  coal,  reacting  with  it,  as  already 


212 


CAEBON  AND   ITS   COMPOUNDS. 


shown,  to  form  carbon  monoxide  and  hydrogen.  The  mixture 
of  these  two  gases  is  then  forced  into  the  "superheaters,"  in 
which  light  petroleum  oils  are  being  decomposed  by  heat. 
Here  the  water-gas  obtains  marsh  gas  and  the  hydrocarbons 
which  give  ordinary  gas  its  illuminating  power. 

From  the  superheaters  the  gas  passes  through  the  "  scrub- 
ber" and  the  "  condenser,"  in  which  the  undecomposed  petro- 
leum and  any  carbon  dioxide,  etc.,  are  removed. 

225.     Comparison  of  the  Two  Kinds  of  Gas.  — The 

composition  of  two  samples  of  illuminating  gas,  one 
made  by  the  distillation  of  coal  and  the  other  by  the 
"  water-gas  "  process,  is  given  in  the  following  table  :  — 


OLD  PROCESS. 

NEW  PROCESS. 

Hydrogen. 

46.0* 

29.5* 

Methane. 

39.0* 

20.0* 

Carbon  monoxide. 

5.5% 

32.7* 

Hydrocarbons. 

6.0* 

11.2* 

Carbon  dioxide. 

1.9* 

2.8* 

Oxygen. 

0.3* 

O.O* 

Nitrogen  (by  difference). 

2.0* 

3.8* 

Total  

100.0$ 

100.0* 

Gas  made  by  the  "water-gas  "  process  is  more  poison- 
ous than  that  made  by  the  distillation  of  coal,  owing  to 


EXERCISES.  213 

the  presence  of  the  large  quantity  of  carbon  monoxide 
in  water-gas ;  otherwise  the  two  gases  do  not  differ 
greatly  in  properties. 

226.  Amount    of    Gas    Used.  —  The    quantity   of 
manufactured  gas   used  in  the  United   States  is  enor- 
mous ;  it  amounts  to  about  sixty  thousand  millions  of 
cubic  feet  annually. 

The  city  of  Chicago  alone  used,  during  the  year  1900, 
not  less  than  seventy-five  hundred  millions  of  cubic 
feet  of  artificial  gas,  to  say  nothing  of  natural  gas. 

227.  Exercises. 

1.  How   many  grams  of   carbon  are   needed  to  reduce  15 
grams  cupric  oxide,  CuO,  if  the  carbon  is  oxidized  to  carbon 
dioxide  ?     How  many  grams  of  carbon  dioxide  are  formed  ? 

2.  What  volume  of  carbon  dioxide  at  0°  C.  and  760  mm.  can 
be  produced  from  840  grams  magnesium  carbonate,  MgCO3? 

3.  How  many  grams  of  sodium  bicarbonate  are  needed  to 
give,  when  heated,  36  liters  of  carbon  dioxide?     (Cf.  §210.) 

4.  What  is   the   formula   of   barium   hydrogen   carbonate  ? 
What  products  are  formed  when  its  solution  is  boiled? 

5.  How  many  liters  of  carbon  dioxide  are  formed,  at  standard 
conditions,  by  the  combustion  of  250  grams  carbon  monoxide  ? 

6.  Calculate  the  percentage  composition  of  calcium  carbon- 
ate, carbon  dioxide,  oxalic  acid. 

7.  How  would  you  distinguish  between  sodium  carbonate, 
sodium  sulphite,  and  sodium  sulphide,  chemically? 

8.  How  would  you  distinguish  between   the   gases  carbon 
dioxide,  carbon  monoxide,  oxygen,  hydrogen,  and  ammonia? 


CHAPTER   XV. 

FLAMES.     HEAT  OF  FORMATION  AND   DECOMPOSITION. 
A.     Flames. 

228.  Luminosity  of  Flames.  —  As  stated  in  Chapter 
II,  §  33,  aflame  is  a  burning  gaseous  body.  The  amount 
of  light  given  off  by  a  flame  depends  upon  the  nature 
of  the  burning  substance  and  upon  its  density. 

As  an  illustration  of  the  influence  of  density,  we  may  take  the 
case  of  hydrogen.  This  substance  ordinarily  burns  in  air  and 
in  oxygen  with,  an  almost  invisible  flame;  if,  however,  the 
hydrogen  and  the  oxygen  are  very  much  compressed  before 
ignition,  the  flame  produced  by  their  union  is  a  very  brilliant 
one. 

The  illuminating  power  of  all  ordinary  flames  is  due 
to  the  presence  of  incandescent  solid  particles.  This 
may  be  illustrated  by  introducing  any  fine  dust  into  a 
flame  of  hydrogen  or  into  the  colorless  Bunsen  flame  ; 
the  flame  at  once  becomes  luminous. 

In  the  combustion  of  substances  containing  carbon,  — 
such  substances  as  candles,  illuminating  gas, -paper,  pe- 
troleum, wood,  coal,  etc., — the  luminosity  of  the  flame 
is  due  to  the  glowing  of  particles  of  carbon  in  the  flame. 
A  cold  object  inserted  into  the  flame  produced  by  one 
of  these  substances  becomes  covered  with  soot ;  and  too 

214 


STRUCTURE   OF  FLAMES. 


215 


little  air,  or  too  much  air,  causes  the  flame  to  smoke, 
owing  to  the  escape  o*f  unburned  particles  of  carbon. 

229.  Structure  of  Flames. -- A  burning  candle 
shows  practically  the  same  phenomena  as  the  other  com- 
pounds of  carbon  just  named,  and  may  be  taken  as  rep- 
resentative of  them. 

The  burning  of  a  small  portion  of  the  wick  of  the 
candle  furnishes  heat  enough  to  melt  some  of  the  wax. 
The  wick  then  draws  the  melted  wax,  by  capillary  action, 
into  the  flame,  where  the  wax  is  first  vaporized  and  then 
ignited. 

If  the  candle  flame  be  examined,  it  will  be  found  to 
consist  of  several  regions,  or  zones,  of  com- 
bustion, surrounding   a    central    cone-shaped       A__-__JJ> 
region  of  unburned  gases.     These  parts  are 

shown  in  vertical  section  in  Fig.  51. 

-B 

X  is  the  region  of  unburned  gases. 

B  is  the  luminous  zone.  It  contains  solid  par- 
ticles of  carbon  in  a  state  of  combustion.  B'  is 
the  ruddy  tip  of  the  luminous  zone. 

A  is  the  outer  mantle  of  the  flame.  Being  non- 
luminous  it  is  obscured  by  the  light  of  B,  except 
at  the  bottom,  where  it  forms  a  blue,  cup-shaped 
region. 

In  addition  to  the  parts  just  named,  an  impor- 
tant region  is  believed  to  exist  about  the  region  X. 
This  zone  is  designated  C  in  the  idealized  section 
of    a  candle  flame  (Fig.   52).      Being    non-luminous,  the   re- 
gion C  is  obscured  by  the  light  of  B. 


— -A 


FIG.  51. 


216 


FLAMES. 


Whether  a  flame  is  to  be  luminous  or  non-luminous  depends 
upon  the  condition  of  affairs  in  the  region  O. 

What  takes  place  in  a  candle  flame  is    probably  as 
follows  :  — 

The  vaporized  paraffin  (wax)  of  the  region  X  (Fig. 
52)  burns  in  part  in  the  zone  (7,  producing 
enough  heat  to  decompose  some  of  the  paraf- 
fin vapor  into  hydrogen,  certain  hydrocar- 
bons (especially  acetylene),  and  solid  carbon. 
These  substances  burn  further  in  the  region 
B,  the  carbon  burning,  as  usual,  with  a  bright 
-X  glow,  and  thus  causing  the  luminosity  of  this 
-—A  regi°n'  I1'1  A  the  gases  and  the  carbon  escap- 
ing unburned  through  B  are  more  or  less 
completely  burned. 

FIG.  52.  230.  Non-Luminous    Flames.  —  The    de- 

composition of  the  combustible  material  of 
region  X  (Fig.  52)  into  acetylene,  carbon,  etc.,  in  the 
region  C  requires  a  definite  degree  of  temperature  ;  hence, 
if  the  temperature  of  0  is  sufficiently  lowered,  the 
flame  becomes  non-luminous.  This  is  exactly  what 
takes  place  in  practice ;  for  if  large  quantities  of  a  cold 
diluting  gas,  e.  g.,  air,  carbon  dioxide,  or  nitrogen,  are 
introduced  into  X,  thus  cooling  C,  the  luminosity  of 
the  flame  is  destroyed.  If,  however,  the  diluting  gases 
are  first  heated,  the  non-luminous  flame  becomes  lumi- 
nous. 


SIMPLE  AND    COMPLEX  FLAMES. 


217 


In  the  non-luminous  flame  the  region  C  may  readily  be 
distinguished  by  its  lightrblue  color. 

The  explanation  of  non-luminosity  in 
flames,  just'  given,  contains  the  theory  of 
the  Bunsen  burner  (Fig.  53).  When  the 
holes-  at  the  base  of  the  burner  are  open, 
the  gas  which  rushes  past  the  holes  draws 
in  currents  of  air,  and  the  flame  is  non- 
luminous;  when  the  holes  are  closed  the 
flame  is  luminous.  Carbon  dioxide,  nitro- 
gen, etc.,  give  the  same  result  as  air. 


Air  «    „„    I  Air 


The  zones  of  combustion  in  the 
Bunsen  flame  are  as  follows  (Fig. 
54):- 


FIG.  53. 


.XT  is  the  region  of  unburned  gas,  as  in  the  candle 
flame. 

G  is  the  light-blue,  inner  cone  surrounding  X. 

B  is  the  non-luminous,  dark  cone.     In  the  candle 
flame  this  is  luminous. 

A.       A  is  the  purple,  outer  mantle.     In  the  candle  this 
is  the  faintly  luminous  liuJlo  surrounding  the  flame. 

231.  Simple  and   Complex  Flames. — The 

C  simplest  flames  are  those  having  only  one  cone 
„  of  combustion.     Illustrations  are  the  flames  of 

.A. 

carbon  monoxide    (<?f.  §  214)  and  hydrogen  (of. 

§  11).  It  is  to  be  noted  that  these  are  sub- 
FIG.  54.  stances  that  can  have  only  one  combustion  in 

air ;  for  carbon  monoxide  can  burn  only  to  form 
carbon  dioxide,  and  hydrogen  only  to  form  water. 


218 


FLAME  b. 


Substances  that  can  have  two  combustions  burn  with 
flames  consisting  of  two  zones. 

The  best  illustration  of  this  fact  would  probably  be  gaseous 
carbon,  if  it  could  be  obtained.  In  place  of  carbon  we  can  use 
cyanogen,  C2N2  (c/.  §  215)  ;  for  the  nitrogen  takes  practically 
no  part  in  the  combustion.  Cyanogen  really  burns  with  aflame 
consisting  of  two  zones.  Both  carbon  and  carbon  monoxide 
burn  to  carbon  dioxide  in  its  flame. 

The  fla'me  of  a  hydrocarbon,  in  which  three  combus- 
tions are  possible,  has,  like  that  of  a  candle,  three 
regions  of  combustion. 

232.  Dissection  of  Flames. - 
Complex  flames  may  be  dissected 
and  the  separate  zones  made  to 
burn  by  themselves.  This  may  be 
done  with  illuminating  gas  as  fol- 
lows :  — 

Two    pieces   of    glass    tubing,    one 
larger  in  diameter  than  the  other  (Fig. 
55),   are   joined  by  a  rubber   tube  so 
that   the   inner   tube   may   be    moved 
within    the  outer  one.     At  first  the 
Gas      inner   tube    is    so  adjusted    that    its 
top   is  just  below  that  of  the  outer 
tube.     The  inner  tube  is  connected 
with   an   inverted  T-tube. 

Illuminating  gas  is  passed  through  one  arm  of  the  T-tube 
and  lighted  at  the  top  of  the  larger  (outer)  tube  ;  then  a  cur- 
rent of  air  is  forced  through  the  other  arm  of  the  T.  By 


FIG.  55. 


ENERGY  CHANGES.  219 

a  careful  adjustment  of  the  air  supply  to  that  of  the  gas,  the 
colorless  Bunsen  flame  is  obtained. 

If,  now,  while  tlie  colorless  flame  is  burning,  the  inner  tube 
is  lowered,  so*  that  the  distance  between  the  tops  of  the  two 
tubes  is  gradually  increased,  a  position  will  be  found  at  which 
the  flame  divides  into  two  flames,  one  upon  the  outer  tube  and 
one  upon  the  inner  tube. 

B.     Heat  of  Formation  and  of  Decomposition. 

233.  Energy  Changes  Accompany  Chemical 
Changes.  —  Chemical  changes  result  not  only  in  the 
formation  of  new  substances,  but  also  in  the  evolution 
or  the  absorption  of  energy.  The  energy  may  appear 
in  the  form  of  heat,  light,  electric  effects,  etc. 

To  illustrate  :  The  union  of  carbon  with  oxygen  produces 
carbon  dioxide  ;  but  at  the  same  time  a  considerable  evolution 
of  heat  and  light  takes  place.  Since  the  energy  evolved  comes 
from  the  carbon  and  the  oxygen,  carbon  dioxide  must  possess 
less  energy  than  the  elements  which  produced  it.  So,  also, 
water  has  less  energy  than  the  hydrogen  and  oxygen  from 
which  it  was  formed. 

A  mixture  of  elements  capable  of  uniting  chemically 
with  evolution  of  heat  must  be  looked  upon  as  having 
potential  energy.  In  the  act  of  union  this  energy  is 
partly  given  up  in  the  kinetic,  i.  e.,  active,  form ;  while 
in  the  resulting  compound  there  must  be  less  energy 
than  existed  in  the  elements.  Furthermore,  to  restore 
the  carbon  and  the  oxygen  of  carbon  dioxide  to  the  ele- 
mentary condition,  as  much  energy  must  be  added  to  the 
carbon  dioxide  as  was  evolved  when  the  elements  united. 


220     HEAT  OF  FORMATION  AND  DECOMPOSITION. 

234.  Heat  of  Formation  and  of  Decomposition.  — 

The  quantity  of  heat  evolved  in  many  cases  of  chemical 
action  lias  been  determined  l>y  experiment.  If  AVC  call 
the  amount  of  heat  necessary  to  raise  the  temperature 
of  one  gram  of  water  from  0°  C.  to  1°  C.  a  ymm-centi- 
t/rtide  (g.  c.)  heat  unit,  we  may  write  the  equation  for 
the  union  of  hydrogen  and  oxygen  as  follows  :  — 

H2  -f  O  =  IT,()  -f  08,400  g.  c.  heat  units. 

This  equation  shows  not  only  that  2  grams  of  hydro- 
gen and  16  grams  of  oxygen  unite  to  form  18  grams  of 
water,  but  also  that  by  their  union  68,400  heat  units 
are  liberated. 

Similarly,  for  the  union  of  12  grams  of  carbon  and  32  grams 
of  oxygen  we  may  write, — 

C  +  O2  =  flQ3  -f-  97,000  g.  c.  heat  units. 

For  the  union  of  1  gram  of  hydrogen  and  35.5  grams  of 
chlorine  the  equation  is,— 

H  -j-  Cl  =  HC1  +  22,000  g.  c.  heat  units. 

The  difference  between  the  energy  (calculated  as  heat) 
possessed  by  the  elements  which  united  to  form  a  com- 
pound and  that  possessed  by  the  compound  itself  is 
called  the  heat  of  formation  of  the  compound. 

The  heat  of  decomposition  of  a  compound  is  numeri- 
cally equal  to  the  heat  of  formation ;  that  is  to  say,  the 
quantity  of  heat  necessary  to  separate  a  compound  into 
its  elements  is  just  as  great  as  that  evolved  when  the 
compound  was  formed  from  its  elements. 


HEAT   OF  FORMATION  EVOLVED  IN  STAGES.      221 

235.  Positive  and  Negative  Heat  of  Formation.  - 

The  heat  of  formation  of  water,  of  carbon  dioxide,  and 
of  hydrochloric  acid  is  positive  (-f-)»  heat  being  evolved 
when  these  compounds  are  formed;  many  cases,  how- 
ever, exist  in  which  heat  is  not  evolved,  but  absorbed 
in  the  formation  of  a  compound  from  its  elements.  In 
such  cases  the  heat  of  formation  is  negative  ( — ). 

An  illustration  is  the  case  of  carbon  disulphide,  a  substance 
which  is  produced  (c/.  §  183)  by  passing  sulphur  vapor  over 
hot  charcoal.  Carbon  burns  in  sulphur  with  absorption  of  heat. 
The  quantity  of  heat  rendered  potential  by  the  union  of  12 
grains  of  carbon  and  64  grams  of  sulphur  is  shown  by  the 
equation,— 

C  -f  2  S  =  CS2  —19,600  g.  c.  heat  units. 
Similarly,  hydrogen  and-iodine  unite  with  absorption  of  heat. 
H  -f  I  =  HI  —  6,100  g.  c.  heat  units. 

A  compound  with  a  negative  heat  of  formation  is  in 
a  state  of  tension,  or  of  unstable  equilibrium ;  for  it  is 
possessed  of  more  energy  than  its  constituent  elements. 
When  such  a  compound  is  decomposed,  energy  is 

evolved. 

236.  Heat  of  Formation  Evolved  in  Stages.  — Just 
as  it  makes  no  difference  in  the  total  quantity  of  heat 
given  out  whether  a  given  mass  of  a  combustible  burns 
slowly  or  rapidly  (cf.  §  28),  so  the  total  amount  of  heat 
evolved  (or  absorbed)  in  the  formation  of  a  compound 
is  the  same  whether  the  compound  is  formed  in  one  or 
in  several  stages. 


222       HEAT   OF  FORMATION  AND   DECOMPOSITION. 

Thus,  the  heat  of  formation  of  calcium  carbonate,  CaCO3,  is 
equal  to  the  sum  of  the  heats  of  formation  of  calcium  oxide,  CaO, 
and  of  carbon  dioxide,  CO2,  plua  the  heat  evolved  when  cal- 
cium oxide  and  carbon  dioxide  unite  to  form  calcium  carbon- 
ate, 


CHAPTER   XVI. 
MOLECULES  AND  ATOMS. 

237.  Law    of    Multiple    Proportions.  —  We     have 

already  learned  (<?f.  §  69)  that  chemical  changes  take 
place  between  definite  masses  of  substances,  and  that 
any  given  chemical  compound  always  contains  the  same 
elements  united  in  the  same  proportion.  While  there 
is  no  reason  to  doubt  these  statements,  a  further  fact  is 
also  true,  and  has  already  been  stated,  viz.,  that  the 
same  two  (or  more)  elements  may  unite  in  different 
proportions  to  form  different  compounds. 

Thus,  carbon  and  oxygen  unite  in  the  proportion  of  3  parts 
of  carbon  to  4  of  oxygen  to  form  carbon  mrmoxide,  and  in  the 
proportion  of  3  of  carbon  to  8  of  oxygen  to  form  carbon  dioxide. 

Similarly,  iron  and  sulphur  form  three  distinct  compounds; 
sulphur  and  oxygen,  two  or  more  (cf.  §  190);  nitrogen  and  oxygen, 
fire  (//.  §§  106  to  171);  potassium,  chlorine,  and  oxygen,  four 
(potassium  salts  of  the  chlorine  acids  named  in  §  10f>). 

All  of  the  cases  named  illustrate  the  Law  of  Multi- 
ple Proportions,  which  may  be  stated  in  its  simplest  form 
as  follows :  — 

If  two  elements  form  several  compounds  with  each  other, 
the  different  masses  of  one  element  which  combine  ivith  a 
fixed  mitxs  of  the  other  /><'<()•  <i  simple  ratio  to  one  another. 


224  MOLECULES  AND  ATOMS. 

Thus,  as  stated  above,  the  ratio  between  the  two  quantities 
of  oxygen  combining  with  3  parts  of  carbon  is  4  :  8,  i.  e.^ 
1  :  2.  The  ratio  between  the  two  quantities  of  oxygen  combined 
with  2  parts  of  sulphur  in  sulphur  dioxide  and  sulphur  dioxide 
is  2:3;  that  between  the  five  quantities  of  oxygen  combined 
with  7  parts  of  nitrogen  is  8: 16  :  24  :  32  :40,  i  e.,  1:  2  :  3  :  4  :  5. 

The  law  of  multiple  proportions  was  stated  by  John 
Dal  ton,  in  1804,  from  the  consideration  of  only  a  few 
compounds ;  but  it  has  been  confirmed  by  the  work  of 
the  past  century,  and  is  one  of  the  fundamental  prin- 
ciples of  Chemistry. 

238.  The   Atomic    Hypothesis.  —  We    may   regard 
matter  as  made  up  in  either  of  two  ways:   (1)  as  infi- 
nitely divisible,  or  (2)  as  composed  of  small,  indivisible 
particles.     We  know  that  a  piece  of  gold,  for  example, 
can  be  divided  into  very  small    pieces  ;    and    Ave    may 
imagine  that  the  subdivision  of  the  gold  might  be  con- 
tinued forever.     On  the  other  hand,  we   may  suppose 
that  after  repeated  subdivision    the    particles    of    gold 
become  so  small  that  they  cannot  be  divided  further. 

Both  of  these  views  of  matter  may  be  reasoned  about,  but 
cannot  be  proved.  The  hypothesis  that  matter  is  composed  of 
indivisible  particles  is  generally  known  as  the  atomic  hypothe- 
sis, or  theory.  The  indivisible  particles  are  called  atoms,  from 
the  Greek  atomos,  meaning  "  indivisible."  At  the  present  time 
chemists  usually  hold  the  atomic  hypothesis. 

239.  The  Law  of  Definite  Proportions  Explained  by 
the  Atomic  Theory.  —  Dalton   saw  that   the   idea   of 


EXPLANATION  OF   THE  LAW.  225 

indivisible  particles  was  connected  with  the  laws  of  defi- 
nite and  multiple  proportions.  For,  if  elements  are 
made  np  of  atoms,  all  the  atoms  of  any  one  element  will 
probably  have  the  same  mass ;  while  atoms  of  different 
elements  will  have  different  masses.  When,  therefore, 
two  elements  unite,  the  union  must  take  place  between 
the  atoms  of  these  elements. 

Suppose  one  atom  of  one  element  (let  us  call  it  A)  unites 
with  one  atom  of  another  element  (B),  and  so  on  throughout 
the  whole  mass  of  the  two  elements  ;  it  is  evident  that  if  there 
is  the  same  number  of  atoms  of  each  kind,  none  of  either  kind 
will  remain  uncombined  when  the  action  is  complete. 

Let  us  suppose,  further,  that  the  atoms  of  B  are  twice  as 
heavy  as  the  atoms  of  A.  Then,  if  the  elements  unite  atom  for 
atom,  the  resulting  compound  will  necessarily  contain  the  ele- 
ments in  the  proportion  of  one  part,  by  weight,  of  A  to  two 
parts,  by  weight,  of  B.  Or,  if  we  analyze  the  compound  con- 
sisting of  A  and  B  united  atom  for  atom,  and  find  that  it  contains 
one  part,  by  weight,  of  A  to  two  parts  of  B,  we  must  conclude 
that  the  atom  of  B  is  twice  as  heavy  as  the  atom  of  A. 

In  the  light  of  the  atomic  theory,  therefore,  chemical 
action  must  take  place  between  definite  masses  of  sub- 
stances. The  theory  is  thus  an  explanation  of  the  law 
of  definite  proportions  (ef.  §  69). 

240.  Explanation  of  th*  Law  of  Multiple  Propor- 
tions. —  The  atomic  theory  is  the  explanation  not  only 
of  the  law  of  definite  proportions,  but  also  of  that  of 
multiple  proportions.  For,  if  atoms  are  indivisible,  as 
their  name  indicates,  elements  that  combine  with  one 


226  MOLECULES  AND   ATOMS. 

another  in  more  than  one  proportion  must  do  so  in  some 
way  that  will  not  require  the  dividing  of  an  atom  ;  that 
is  to  say,  the  elements  A  and  B  must  unite  in  the  pro- 
portion of  one  atom  of  A  to  one  of  B,  or  one  of  A  to  two 
of  B,  or  one  of  A  to  three  of  B,  or  two  of  A  to  three  of 
B,  etc. 

Let  us  suppose,  as  in  §  239,  that  the  atom  of  B  is  twice  as 
heavy  as  that  of  A  ;  then,  if  A  and  B  combine  atom  for  atom, 
it  is  evident  that  the  compound  formed  will  contain  the  elements 
in  the  proportion  of  one  part,  by  weight,  of  A  to  two  parts  of  B; 
but  if  one  atom  of  A  unites  with  two  atoms  of  B,  the  resulting 
compound  will  contain  the  elements  in  the  proportion  of  one 
part,  by  weight,  of  A  to  four  parts  of  B. 

Thus,  if  we  assume  the  atomic  hypothesis,  it  follows 
that  elements  which  combine  with  one  another  in  more 
than  one  proportion  must  do  so  according  to  the  law  of 
multiple  proportions. 

241.  Distinction  between  Atoms  and  Molecules.  - 

As  has  already  been  stated  (cf.  §  131),  the  physical 
properties  of  gases  receive  a  reasonable  explanation 
from  the  assumption  that  matter  is  composed  of  mole- 
cules. Molecules  and  atoms  are  not  identical ;  for  while 
atoms  are  thought  of  as  the  very  smallest  particles  into 
which  matter  is  capable  of  being  divided,  molecules  are 
held  to  be  the  aggregations  of  atoms  which  form  the 
physical  units  of  matter.  The  atoms  composing  a 
molecule  do  not  (usually)  part  company  when  matter 
undergoes  a  physical  change. 


MOLECULAR  MASSED  OF  GASEOUS  SUBSTAXCES.  227 

To  illustrate  :  The  physical  properties  of  ammonia  are  de- 
termined by  the  properties  of  the  ammonia  molecules'  it  is  only 
when  we  subject  ammonia  to  a  chemical  change  that  we  divide 
the  molecule.  Then  it  is  that  the  properties  of  the  nitrogen 
and  hydrogen  atoms  come  into  play. 

The  molecules  of  elements  consist  of  atoms  of  only 
one  kind ;  while  the  molecules  of  compound  substances 
contain  atoms  of  two  or  more  kinds. 

242.  The  Molecular  Masses  of  Gaseous  Substances. 

-According  to  Avogadro's  Rule  (</.  §  133),  the  num- 
ber of  molecules  in  equal  volumes  of  all  gaseous  sub- 
stances is  approximately  the  same.  This  means  that  a 
liter  of  chlorine,  for  example,  contains  practically  as 
many  chlorine  molecules  as  there  are  hydrogen  mole- 
cules in  a  liter  of  hydrogen,  or  hydrochloric  acid  mole- 
cules in  a  liter  of  hydrochloric  acid  gas,  if  only  the 
temperature  and  pressure  are  the  same  in  all  cases. 

The  hypothesis  applies  not  only  to  true  gases,  but  also  to  the 
vapors  of  many  substances  ordinarily  liquid  or  solid.  Thus,  a 
liter  of  steam,  or  one  of  acetic  acid  vapor,  is  assumed  to  con- 
tain practically  as  many  molecules  as  a  liter  of  hydrogen  under 
the  same  conditions. 

From  the  relative  masses  of  equal  volumes,  i.  e.,  the 
densities,  of  gaseous  substances,  we  can  get  the  relative 
masses  of  molecules ;  for,  if  a  liter  of  chlorine  is  found 
to  weigh  3.18  grams  and  a  liter  of  nitrogen,  under  the 
same  conditions,  1.25  grams,  and  if  the  number  of 
molecules  in  a  liter  of  each  is  the  same,  the  relative 


228  MOLECULES  AND   ATOMS. 

masses  of  the  chlorine  and  nitrogen  molecules  must  be 
as  3.18  :  1.25. 

Moreover,  since  the  relative  masses  of  equal  volumes 
of  hydrogen,  hydrochloric  acid  gas,  steam,  and  oxygen, 
under  the  same  conditions,  are  as  0.0896  :  1.63  :  0.8064 
:  1.43,  respectively,  these  numbers  must  express  the 
relative  masses  of  the  molecules  of  the  substances 
named. 

Chemists  have  adopted  as  a  standard  the  molecular 
mass  of  oxygen  and  have  called  it  32.  The  reason  for 
this  will  appear  later  (<?/".  §  256). 

The  molecular  mass  of  hydrogen  is  thus  determined 
from  the  proportion,  — 

Wt.  of  1  liter  of  oxygen  :  Wt.  of  1  liter  of  hydrogen  :: 
molecular  mass  of  oxygen  :  molecular  mass  of  hydrogen  ;  or, 

1.43  :  0.0896  ::  32  :  x. 
Whence  x  =  2. 005. 

By  similar  processes  the  molecular  masses  of  chlorine, 
nitrogen,  hydrochloric  acid,  and  steam  may  be  deter- 
mined to  be  approximately  71,  28,  36.5,  and  18,  re- 
spectively. 

243.  Vapor  Density  Methods.  —  It  is  evident  from 
the  preceding  section  that  we  can  determine  the  approxi- 
mate molecular  mass  of  a  gaseous  substance  or  of  a  sub- 
stance that  can  readily  be  vaporized.  We  need  only 
get  the  weight  of  a  given  volume  of  the  substance  in 
the  gaseous  state  and  compare  this  weight  with  that  of 


OTHER   METHODS. 


229 


an  equal  volume  of  oxygen.     Methods  for  determining 
this  ratio  are  called  "  vapor  density  "  methods. 

Suppose  we  wish  to  get  the  approximate  molecular  mass  of 
chloroform,  a  substance  boiling  at  61°  C.,  at  ordinary  pressure. 
Several  vapor  density  methods  are  available;  one  of  them  is 
as  follows  :  — 


Victor  Meyer's  Method.  —  In  the  method  of  Victor  Meyer 
(Fig.  56),  a  weighed  amount  of  the  chloro- 
form (or  other  substance  whose  vapor  density 
is  to  be  found)  is  dropped  into  the  tube  A 
and  vaporized ;  the  air  which  is  expelled 
through  the  delivery  tube  and  collected  in 
the  graduate.d  tube  is  a  measure  of  the  vol- 
ume of  the  chloroform  vapor.  We  thus  ob- 
tain the  weight  of  a  given  number  of  cubic 
centimeters  of  chloroform  vapor.  By  com- 
paring this  weight  with  the  weight  of  an 
equal  volume  of  oxygen,  we  can  get  the 
molecular  mass  of  chloroform. 

The  tube  A  is  raised  to  any  required  tem- 
perature by  boiling  some  liquid  in  the  jacket 
B  ;  in  the  case  of  chloroform,  water  might 
be  used. 

244.  Other  Methods  of  Determining 
Molecular    Masses.  -  -  Vapor     density 
methods      of      determining     molecular 
masses  do  not  apply  to  substances  that 
cannot   be    vaporized    without    decomposition    nor    to 
those    whose    boiling    temperatures    are    so  high   as  to 
be    practically    unattainable.      For    soluble    substances, 


FIG.  56. 


230  MOLECULES  AND  ATOMS. 

however,  other  methods  can  be  used.  These  methods 
are  based  upon  the  fact  that  substances  in  dilute  solu- 
tion behave  much  as  they  would  if  they  were  gaseous 
(cf.  §§138  and  1«39)  and  occupied  the  same  volume  as 
is  occupied  by  the  solution. 

Thus,  if  we  take  two  substances  whose  molecular  masses 
have  been  determined  in  other  ways  and  dissolve  equal  quan- 
tities of  them  in  equal  amounts  of  the  same  solvent,  we  shall 
find  that  the  osmotic  pressures  of  the  two  solutions  are  propor- 
tional to  the  molecular  masses  of  the  dissolved  substances.  That 
is  to  say,  the  number  of  the  dissolved  molecules,  and  not  their 
nature,  determines  the  osmotic  pressure. 

By  comparing  the  osmotic  pressure  of  a  substance  of  known 
molecular  mass  with  that  of  one  whose  molecular'  mass  is  itn- 
known,  the  unknown  molecular  mass  might  be  found. 

The  osmotic  pressure  method  is  not  used  much,  be- 
cause two  other  effects  due  to  dissolved  molecules  are 
more  readily  measured;  these  are  (1)  the  raising  of  the 
boiling  point  of  the  solvent  and  (2)  the  depression  of  its 
freezing  point. 

245.  Boiling  Point  and  Freezing  Point  Methods.  - 
If  we*  compare  the  temperature  of  boiling  pure  water 
with  that  of  water  containing  a  dissolved  substance, 
e.  g.,  sugar,  we  shall  find  that  the  dissolved  substance 
causes  the  water  to  boil  at  a  higher  temperature.  This 
is  true  in  general;  a  dissolved  substance  raises  the  boil- 
ing temperature  of  the  solution  above  that  of  the  pure 
solvent. 


OBTAINING  EXACT  MOLECULAR  MASSES.        231 

For  dilute  solutions,  increasing  the  relative  quantity  of  the 
dissolved  substance  increases  the  rise  of  the  boiling  tempera- 
ture proportionally. 

Moreover,  if  two  substances  of  known  molecular 
masses  are  dissolved  in  amounts  proportional  to  their 
molecular  masses  in  equal  quantities  of  the  same  sol- 
vent, the  rise  of  the  boiling  point  is  the  same  in  the  two 
cases.  But,  taking  substances  in  quantities  proportional 
to  their  molecular  masses  is  the  same  as  taking  equal  num- 
bers of  molecules ;  hence  the  rise  of  the  boiling  point, 
like  the  osmotic  pressure,  is  proportional  to  the  number 
of  dissolved  molecules. 

The  freezing  point  of  a  solution,  on  the  other  hand, 
is  loiver  than  that  of  the  pure  solvent;  but  the  amount 
that  the  freezing  point  is  depressed,  like  the  rise  in  the 
boiling  temperature,  is  proportional  to  the  number  of 
dissolved  molecules,  and  independent  of  their  nature. 
Therefore,  if  substances  are  dissolved  in  equal  amounts 
of  a  given  solvent  in  quantities  proportional  to  their 
molecular  masses,  the  freezing  points  of  the  solutions  are 
lowered  to  the  same  degree. 

There  are  thus  both  freezing  point  and  boiling  point  metk- 
ods  for  determining  the  molecular  masses  of  many  substances 
that  cannot  be  handled  by  vapor  density  methods. 

246.  Methods  of  Obtaining  Exact  Molecular  Masses. 

—  All  of  the  me  chods  described  give  only  approximate 
molecular  masses  ;  exact  molecular  masses  are  found  by 
quantitative  analysis. 


232  MOLECULES  ANT)  ATOMS. 

The  methods  used  may  be  illustrated  by  the  case  of  acetic 
acid,  HC2H3O2 ;  the  exact  molecular  mass  of  this  substance 
may  be  found  by  a  study  of  its  silver  salt,  AgC2H3O2. 

Silver  acetate  contains  silver,  carbon,  hydrogen,  and  oxygen. 
The  per  cent  of  the  silver  being  found  by  analysis  to  be  64.65, 
that  of  the  remainder  of  the  molecule  must  be  35.35.  The 
combining  proportion  of  silver,  if  we  take*the  atomic  mass  of 
oxygen  as  exactly  16,  is  very  nearly  107.94;  the  mass  of  all 
of  the  silver  acetate  molecule  except  the  silver  is,  therefore, 
found  from  the  proportion, 

107.94  :  x  ::   64.65  :  35.35. 
Hence  x  =  59.02. 

Since  acetic  acid  is  silver  acetate  with  the  silver  replaced  by 
hydrogen,  we  must  add  to  59.02  the  number  representing  the 
mass  of  the  hydrogen  that  107.94  parts  of  silver  replace.  This 
is  1.003.  Hence  the  exact  molecular  mass  of  acetic  acid  is 
60.023. 

247.  Atomic  Masses. --The  simplest  conception 
that  has  been  formed  regarding  the  constitution  of  mat- 
ter is  Dalton's  atomic  theory  (cf.  §  238).  An  atom  is 
assumed  to  have  a  real  and  fixed  mass,  just  as  truly  as 
a  pound  of  iron  has ;  this  mass  is,  however,  so  small 
that  we  have  not  the  ability  to  determine  it.  Chemists 
are,  therefore,  concerned  only  about  the  relative  masses 
of  the  different  kinds  of  atoms. 

Relative  atomic  masses  imply  a  standard ;  the  stand- 
ard atomic  mass,  that  of  oxygen,  is  assumed  to  be  ex- 
actly 16  (cf.  §  256). 

When  we  say  that  chlorine  has  an  atomic  mass  of  35.45, 
we  mean  that  the  relation  between  the  mass  of  the  chlorine 


OF  ATOMIC  MASSES. 


233 


atom  and  that  of  the  oxygen  atom  is  as  35.45  :  16.  The  hy- 
drogen atom  has  a  relative  atomic  mass  of  1.003.  We  usually 
give  atomic  masses  in  round  numbers.  Thus,  the  atomic 
mass  of  chlorine  is  said  to  be  35.5,  and  that  of  hydrogen,  1. 

248.  Determination  of  Atomic  Masses.  —  The  selec- 
tion of  the  atomic  mass  of  an  element  is  no  easy  matter, 
especially  if  the  element  has  only  a  few  known  com- 
pounds. 

How  the  atomic  mass  of  chlorine  comes  to  be  about 
85.5,  is  seen  from  the  accompanying  table  of  compounds 
of  chlorine. 


MASS  OF 

COMPOUNDS 

MOLECULAR 

PARTS 

CHLORINE  IN 

OF 

MASS  OF 

PER  CENT  OF 

THE  MOLECULE 

CHLORINE. 

COMPOUND. 

CHLORINE. 

OF 

COMPOUND. 

Hvdrogen  chloride. 

3G.5 

97.26 

35.5 

Acctyl  chloride. 

78.5 

45.22 

35.5 

Ethyl  chloride. 

64.5 

55.03 

35.5 

Carbonyl  chloride. 

99. 

71.71 

71. 

Chromium  oxychloride. 

155.1 

45.8 

71. 

Chlorine  gas. 

71. 

100. 

71. 

Arsenic  trichloride. 

181.5 

58.67 

106.5 

Phosphorus  trichloi'ide. 

137.5 

77.45 

106.5 

Carbon  tetrachloride. 

154. 

92.21 

142. 

Silicon  tctrachloride. 

170.4 

83.33 

142. 

Tantalic  chloride. 

3<i0.5 

49.24 

177.5 

ITcxachlorethane. 

.     237. 

89.87 

213. 

The  molecular  mass  of  each  of  the  above  compounds 
can  be  determined  by  vapor  density  methods  (cf.  §  243). 
These  compounds  having  been  analyzed,  the  per  cent  of 
each  constituent  is  known. 


234  MOLECULES  AND   ATOMS. 

Thus,  the  per  cent  of  chlorine  in  hydrochloric  acid  is  97.20. 
This  number,  must,  therefore,  be  to  100,  the  total  per  cent,  as 
the  mass  of  chlorine  in  the  molecule  is  to  the  mass  of  the 
ivhole  molecule,  or, 

97.26  :  100  ::  x  :  36.5. 
Whence  x  =  35.5,  approximately. 

la  the  same  way  the  mass  of  all  the  chlorine  atoms  contained 
in'the  molecule  of  each  of  the  other  compounds  can  be  found. 

It  will  be  observed  that  the  greatest  common  factor  of 
all  the  numbers  representing  the  masses  of  chlorine  in 
the  molecules  of  the  compounds  given  in  the  list  is  35.5  ; 
a  study  of  the  other  compounds  of  chlorine  only  confirms 
the  choice  of  this  number.  Thirty-five  and  five-tenth*  is, 
therefore,  taken  as  the  approximate  atomic  mass  of 
chlorine. 

The  general  method  of  selecting  the  atomic  masses  of 
such  elements  as  hydrogen,  nitrogen,  bromine,  iodine, 
carbon,  sulphur,  phosphorus,  arsenic,  and  some  others  is 
as  given  for  chlorine.  We  get  the  molecular  mass  and 
the  percentage  composition  of  each  of  the  known  com- 
pounds of  the  element ;  the  greatest  common  factor  of  the 
numbers  representing  the  masses  of  the  element  found 
in  the  molecules  of  all  its  compounds  will  be  the  atomic 
mass  of  the  element. 

Note  that  the  atomic  mass  found  in  this  way  is  the  greatest 
atomic  mass  the  element  can  have.  The  study  of  other  com- 
pounds of  the  element  may  make  it  necessary  for  chemists  to 
choose  some  sub-multiple  of  the  maximum  atomic  mass,  but 
never  a  multiple  of  it. 


DULONG   AND   PETIT' S  RULE.  235 

249.  Exact    Atomic    Masses.  —  If     the    molecular 
masses  of  the  compounds  from  a  consideration  of  which 
the  atomic  mass  of  an  element  is  chosen  have  been  deter- 
mined   accurately,  i.  e.,  by  quantitative    methods,    the 
atomic  mass  of  the  element  will  be  accurate  ;  but  if  the 
molecular  masses  of  the  compounds  were  determined  by 
vapor  density  methods,  as  in  the  case   of  the  chlorine 
compounds    in   §  248,  the    atomic    mass    will    be    only 
approximate. 

Exact  atomic  masses  are  determined  by  comparison 
with  oxygen,  whose  atomic  mass  is  chosen  16.  Formerly 
hydrogen  was  used  as  the  standard,  its  atomic  mass  be- 
ing chosen  1.  But  oxygen  forms  compounds  with  so 
many  more  elements  than  hydrogen  does,  that  atomic 
masses  can  be  determined  more  directly,  and  hence 
more  accurately,  if  oxygen  is  taken  as  16. 

The  list  of  atomic  masses  is  given  in  the  Appendix. 

250.  Dulong  and  Petit's  Rule.  —  A  rapid  method  of 
getting  at  the  approximate  atomic  masses  of  solid  ele- 
ments, especially  'metals,  is  based  upon  the  relation  be- 
tween the  specific  heat  (better,  relative  thermal  capacity) 
of  an  element  and  its  atomic  mass. 

By  the  specific  heat  of  a  substance  ice  mean  the  quantity  oj 
heat  required  to  raise  the  temperature  of  a  certain  mass  of  the 
substance  one  degree  of  temperature  as  compared  with  the  amount 
needed  to  raise  the  temperature  of  an  equal  mass  of  water  one 
degree.  To  illustrate  :  if  two  iron  balls  of  equal  mass  and  at 
the  same  temperature,  say  at  200°  C.,  are  put  into  equal  and 


236  MOLECULES  AND  ATOMS. 

sufficiently  large  masses  of  water  and  mercury,  it  is  evident 
that  the  same  quantity  of  heat  will  be  imparted  in  the  two  cases. 
The  temperature  effect  will,  however,  be  very  different ;  foi 
the  mercury  will  be  heated  through  about  32  times  as  man^ 
degrees  as  the  water. 

Thus  the  specific  heat  of  mercury  is  aV,  or  O.OH19. 

In  1819,  Dulong  and  Petit  observed  the  existence  oi 
the  rule  which  is  called  by  their  names ;  the  rule  may 
be  stated  as  follows  :  — 

The  specific  heat  of  a  solid  element  multiplied  by  its 
atomic  mass  is  a  constant  (about  6.25).  Some  illustra- 
tions appear  in  the  following  table :  — 


ELEMENT. 

SPECIFIC  HEAT. 

ATOMIC  MASS. 

ATOMIC  HEAT. 

So<liutn. 

0.29 

X 

23 

_ 

6.7 

Potassium. 

0.166 

X 

39 

= 

6.5 

Iron. 

0.112 

X 

56 

= 

6.3 

Silver. 

0.057 

X 

108 

= 

6.1 

Tin. 

0.054 

X 

118 

— 

6.4 

Gold. 

0.032 

X 

196 

= 

6.3 

Mercury  (solid).         , 

0032 

X 

200 

= 

6.4 

Lead. 

0.031 

X 

206.4 

= 

6.4 

The  table  shows  that  the  higher  the  atomic  mass  is,  the 
lower  is  the  specific  heat.  The  complete  list  of  specific  heats 
is  given  in  the  Appendix. 

It  is  apparent  that  the  rule  of  Dulong  and  Petit  may 
be  used  to  determine  the  approximate  atomic  mass  of  an 
element.  Thus,  knowing  the  specific  heat  of  cadmium 


APPLICATION   OF  ATOMIC  MASS  METHODS.       237 

to  be  0.0567,  we  can  at  once  get  its  approximate  atomic 
mass.     We  simply  solve  for  x  in  the  equation, 


6.25 

=  x. 


0.0567 
Hence  x  =  110. 

The  exact  atomic  mass  of  cadmium  is  about  112. 

251.  Application  of  Atomic  Mass  Methods.  —  The 

methods  used  in  actually  getting  the  atomic  mass  of  an 
element  may  be  illustrated  by  the  the  case  of  zinc.  The 
method  used  for  chlorine  will  not  apply  here  ;  we  must, 
therefore,  study  the  action  of  zinc  with  elements  of 
known  atomic  mass,  e.  g.,  with  hydrogen,  chlorine,  and 
oxygen. 

When  zinc  is  dissolved  in  certain  dilute  acids,  it  sets  free 
hydrogen;  we  can  thus  get  the  relation  between  the  mass  of 
zinc  taken  and  that  of  the  hydrogen  displaced,  i.  e.,  the  equiva- 
lent of  the  zinc.  Now,  if  for  each  atom  of  zinc  dissolved  an 
atom  of  hydrogen  is  set  free,  the  mass  of  the  zinc  must  be  to 
the  mass  of  the  hydrogen  as  the  atomic  mass  of  zinc  is  to  the 
atomic  mass  of  hydrogen.  Since  we  know  the  atomic  mass  of 
the  hydrogen,  we  could  calculate  that  of  the  zinc. 

The  relation  of  zinc  to  hydrogen  is  about  as  32.7  :  1. 

A  second  element  with  which  we  can  compare  zinc  is 
chlorine.  The  compound  of  these  two  elements,  zinc  chloride, 
may  be  made  either  (1)  by  dissolving  zinc  in  hydrochloric 
acid  and  evaporating  the  solution,  or  (2)  by  burning  zinc  in 
chlorine, 


238  MOLECULES  AND   ATOMS. 

The  relation  of  zinc  to   chlorine  in  zinc  chloride,  as 
determined  by  quantitative  analysis,  is  about  as  32.7  : 
35.5.     Hence,  if  chlorine  and  zinc  have  united  atom  for 
atom,  and  the  atomic  mass  of  chlorine  is  35.5,  the  atomic 
mass  of  zinc  must  be  32.7. 

Let  us,  however,  consider  a  third  compound  of  zinc, 
viz.,  its  oxide. 

This  substance  may  be  made  by  heating  zinc  in  oxygen,  or, 
better,  by  dissolving  zinc  in  dilute  nitric  acid  and  heating  the 
zinc  nitrate  formed  to  a  high  temperature. 

The  proportions  of  the  elements  in  zinc  oxide  are,  — 
zinc,  65.4  parts,  to  oxygen,  16  parts. 

Here,  also,  we  might  assume  that  zinc  and  the  other 
element  are  united  atom  for  atom ;  if  so,  the  atomic 
mass  of  zinc  is  65.4' 

To  compare  these  results  let  us  write  them  down  together  :  — 

(1)  One  part  of  hydrogen  (1  atom)  was  replaced  by  32.7 
parts  of  zinc. 

(2)  Thirty- five  and  five-tenths  parts  of   chlorine  (1   atom) 
united  with  32.7  parts  of  zinc. 

(3)  Sixteen  parts  of  oxygen  (1  atom)  united  with  65.4  parts 
of  zinc. 

Evidently  we  cannot  assume  that  one  atom  of  zinc 
unites  with  one  atom  of  the  other  element  in  each  of 
the  three  cases  ;  for  we  cannot  have  some  zinc  atoms 
with  a  mass  of  32.7  each,  and  others  with  a  mass  twice 
as  great. 

If  the  atomic  mass  of  zinc  is  32.7,  one  atom  of  oxygen  must 


APPLICATION  OF  ATOMIC  MASS  METHODS.      239 

unite  with  two  atoms  of  zinc  to  form  zinc  oxide  ;  if,  however, 
the  atomic  mass,  of  zinc  is  65.4,  two  atoms  of  hydrogen  must 
have  been  replaced  by  one  of  zinc,  and  two  atoms  of  chlorine 
must  have  united  with  one  of  zinc. 

The  molecular  mass  oi  zinc  chloride,  or  of  zinc  oxide, 
would  help  us ;  that  of  zinc  chloride  is  the  one  used ; 
for  the  weight  of  a  known  volume  of  its  vapor  has  been 
obtained.  By  a  solution  of  the  proportion, 

Wt.  of  1  liter  of  zinc  chloride  vapor  :  wt.  of  1  liter  of  oxygen 
(under  the  same  conditions)  : :  molecular  mass  of  zinc  chloride  * 
molecular  mass  of  oxygen,  i.  e.,  32  (c/.  §  242), 

the  molecular  mass  of  zinc  chloride  was  found  in  an 
actual  experiment  to  be  about  134.  This  result  is  suf- 
ficiently accurate  to  enable  us  to  decide  that  the  molec- 
ular mass  of  zinc  chloride  is  136.4  rather  than  one-half 
or  twice  this  number. 

A  molecule  of  zinc  chloride  thus  contains  65.4  parts 
of  zinc  and  71  parts  (two  atoms)  of  chlorine.  We  do 
not,  however,  know  whether  the  two  atoms  of  chlorine 
are  united  with  one  atom  of  zinc  having  a  mass  of  65.4, 
or  with  two,  each  having  a  mass  of  32.7  ;  or  with  three, 
each  having  a  mass  of  21.8,  etc. 

We  now  apply  Dulong  and  Pe tit's  rule.  By  substi- 
tuting the  specific  heat  of  zinc,  found  by  experiment 
(0.094),  in  the  expression, 

6.25 


specific  heat 


=  atomic  mass, 


240  MOLECULES  AND  ATOMS. 

we  obtain  66.5  for  the  approximate  atomic  mass  of  zinc. 
This  shows  us  that  the  number  65.4  is  to  be  taken 
rather  than  any  sub-multiple  of  it. 

252.  How  Formulas  are  Determined.  — If  we  know 
the  molecular  mass  of  a  compound,  the  parts  per  cent  of 
each  element  in  the  compound,  and  the  atomic  masses  of 
the  elements,  we  can  determine  the  formula  of  the 
compound. 

To  illustrate:  A  certain  compound  of  carbon  and  hydrogen 
is  found  by  vapor  density  methods  to  have  the  molecular  mass 
72.  The  per  cent  of  carbon  is  83.33  and  that  of  the  hydrogen 
16.67.  The  mass  of  all  the  carbon  atoms  in  the  molecule 
must  be  60  (=72  X  .8333),  and  the  mass  of  all  the  hydrogen 
atoms  12.  Since  the  mass  of  each  carbon  atom  is  12,  there 
must  be  5  atoms  of  carbon  in  the  molecule  of  the  compound  ; 
and  since  the  mass  of  each  hydrogen  atom  is  1,  there  must  be 
12  hydrogen  atoms.  Hence  the  formula  is  C5H12. 

If  the  molecular  mass  is  not  known,  we  cannot  be 
sure  whether  we  have  the  correct  formula  or  not.  Thus 
the  formulas  C10H24,  C20H48,  etc.,  would  all  have  the 
same  quantitative  composition  as  C5H12.  This  may  be 
illustrated  .by  the  following  case  :  — 

The  substances  formaldehyde,  acetic  acid,  and  grape- 
sugar  all  have  the  same  quantitative  composition,  viz. :  — • 

Carbon,  40.00%; 
Hydrogen,  6.67% ; 
Oxygen,  53.33%. 

What  are  their  formulas  ? 


MOLECULAR  FORMULAS  AND   EQUATIONS.      241 

We  can  get  the  relation  between  the  numbers  of  atoms  of 
each  element  in  the  molecule  by  dividing  the  per  cent  of  each 
element  by  the  atomic  mass.  This  gives, — 

C,  3.33  (=40-4-12); 
H,  6.67  (=6.67  -4-1); 
O,  3.33  (=53.33-4-16). 

Since  there  cannot  be  3.33  or  6.67  atoms,  we  rind  the  simplest 
relation  between  these  numbers. 

3.33  :  6.67:  3.33  ::  l :  2:  l. 

Hence,  the  simplest  formula  a  substance  of  the  given 
composition  could  have  is  CH2O.  Such  a  substance 
would  have  a  molecular  mass  of  30.  This  is  the  mo- 
lecular mass  of  formaldehyde;  hence,  the  formula  of 
formaldehyde  is  CH2O.  The  molecular  mass  of  acetic 
acid  is  60,  and  that  of  grape-sugar,  180;  hence  the 
formulas  of  these  substances  are  C9H,O9  and  CRH10(X 

&       4      &  o       12      b 

respectively. 

253.  Molecular   Formulas    and   Equations.  -  -  The 

symbol  of  an  element  means  not  only  the  element  in  gen- 
eral (cf.  §  5)  and  a  definite  mass  of  it,  but  also  an  atom 
of  it.  As  symbols  stand  for  atoms,  so  formulas  represent 
molecules.  The  numbers  which  we  called  combining  pro- 
portions in  §  75  and  §  76,  because  they  represent  the 
proportions  by  weight  in  which  elements  and  compounds 
enter  into  reactions,  were  derived  from  the  commonly 
accepted  atomic  and  molecular  masses.  The  equations 
previously  given  may  thus  be  looked  upon  as  represent- 


242  MOLECULES  AND  ATOMS. 

ing  changes  in  the  arrangement  of  the  atoms  composing 
molecules. 

The  equation  for  the  action  of  zinc  and  sulphuric  acid,  fo 
example,  means  that  one  atom  of  zinc  reacts  with  one  molecul 
of  sulphuric  acid  to  form  a  molecule  of  zinc  sulphate  and  tw< 
atoms  of  hydrogen.  The  hydrogen  we  obtain  by  this  reactioi 
is  not,  however,  hydrogen  in  the  form  of  atoms,  but  of  mole 
cules  ;  for  the  atoms  have  united  with  one  another,  every  tw< 
forming  a  molecule  of  hydrogen.  Hence  we  write  H2,  meaning 
one  molecule  of  hydrogen  (consisting  of  two  atoms),  rather  thai 
2  H,  which  means  simply  two  atoms  of  hydrogen.  For  the  sann 
reason  we  write  the  equation  for  the  union  of  hydrogen  an< 
oxygen,  2  H2  -f  O2  =  2  H2O;  and  not  4 II  +  2  O  =  2 II2O;  no 
yet  2  H  +  O  =  II2O. 

254.  Nascent,  or  Atomic,  State.  —  In  general,  w< 
represent  the  factors  and  the  products  of  equations  b 
molecular  formulas  instead  of  by  atomic  symbols.  W« 
do  this  because  we  believe  that  elements  in  the  fre 
condition  usually  consist  not  of  atoms  but  of  molecule 
made  up  of  two  or  more  atoms  united. 

Thus,  although  oxygen,  hydrogen,  chlorine,  etc.,  are  se 
free  from  the  molecules  of  their  compounds  in  the  form  o 
atoms,  they  do  not  remain  so;  for  the  atoms  unite  to  forn 
molecules.  In  fact,  an  element  like  oxygen  is  in  the  atomi 
condition  only  during  the  infinitely  short  time  that  intervene 
between  the  liberation  of  the  atoms  from  molecules  and  thei 
union  with  other  atoms  to  form  new  molecules. 

A  proof  of  this  is  the  fact  that  an  element  has  proper 
ties  at  the  instant  of  its  liberation  from  a  compound 


NUMBER    OF  ATOMS  IN  MOLECULES.  243 

i.  e.,  in  its  nascent  condition,  which  differ  from  those 
we  ordinarily  recognize  as  belonging  to  the  element. 
Thus,  when  chlorine  liberates  oxygen  from  water  (ef. 
§  88)  in  the  process  of  bleaching,  the  nascent  oxygen  is 
much  more  active  than  ordinary  oxygen.  Many  other 
cases  of  the  same  kind  are  known. 

The  nascent  state  and  the  atomic  state  of  an  element 
thus  coincide. 

255.  Number  of  Atoms  in  the  Molecules  of  Ele- 
ments. —  We  have  just  learned  that  the  molecules  of 
hydrogen  and  oxygen  consist  of  two  atoms  each;  the 
same  fact  is  true  of  several  other  elementary  gases. 

We  can  reason  that  the  molecules  of  oxygen  must 
consist  of  at  least  two  atoms  each  from  the  proportions 
by  volume  in  whicli  hydrogen  and  oxygen  combine. 
We  can  represent  these  proportions  graphically  by  let- 
ting squares  represent  equal  volumes  of  the  gases  hydro- 
gen and  oxygen,  and  of  the  steam  formed  by  their  union. 

CD        +        D  OH 

2  vols.  hydrogen     1  vol.  oxygen     2  vols.  steam. 

Now,  according  to  Avogadro's  rule,  the  number  of 
molecules  of  hydrogen  and  oxygen  in  equal  volumes 
must  be  approximately  equal ;  hence,  — 

Two  molecules  of  hydrogen  -j-  1  molecule  of  oxygen  give  2 
molecules  of  steam. 

To  form  one  molecule  of  steam,  however,  we  must 
imagine  union  to  take  place  between  one  molecule  of 


244  MOLECULES  AND  ATOMS. 

hydrogen  and  one-half  a  molecule  of  oxygen.     This  half- 
molecule  of  oxygen  is  the  atom  of  oxygen. 

That  the  molecules  of  hydrogen  and  of  chlorine  contain  two 
atoms  is  evident  from  a  similar  method  of  reasoning;  for,  since 
one  volume  of  hydrogen  and  one  volume  of  chlorine  give  two 
volumes  of  hydrochloric  acid  (c/.  §  94),  one  molecule  of  hydro- 
gen and  one  molecule  of  chlorine  must  give  two  molecules  of 
hydrochloric  acid.  Therefore,  one  molecule  of  hydrochloric 
acid  must  contain  one-half  a  molecule  of  hydrogen  and  one- 
half  a  molecule  of  chlorine.  Here  again  the  half-molecules  are 
the  atoms  of  these  gases. 

Methods  are  known  for  determining  the  number  of 
atoms  in  the  molecule  of  an  element  in  the  gaseous 
state  ;  fche  results  obtained  agree  with  those  deduced 
by  our  reasoning.  No  reactions  now  known  make  it 
necessary  for  us  to  assume  that  there  are  more  than  two 
atoms  in  the  molecules  of  hydrogen,  oxygen,  nitrogen, 
chlorine,  and  fluorine.  Some  other  elements,  not  ordi- 
narily gaseous,  have  two  atoms  to  the  molecule  in  the 
gaseous  state.  In  other  cases  the  molecule  consists  of 
only  one  atom ;  in  others  still,  of  three  or  more  atoms. 

256.  Reason  for  Choosing  32  as  the  Molecular 
Mass  of  Oxygen.  —  We  can  see  now  why  the  molecular 
mass  of  oxygen  was  taken  as  32  (<?/'.  §  242).  For,  if 
the  mass  of  the  oxygen  atom,  i.  e.,  the  half -molecule,  is 
assumed  to  be  exactly  16  (cf.  §  249),  the  mass  of  the 
molecule  must  be  twice  as  great.  Formerly,  when  the 
mass  of  the  hydrogen  atom  was  taken  as  the  standard 


LAW  OF  SIMPLE    VOLUMES.  245 

for  atomic  masses,  the  mass  of  the  hydrogen  molecule 
(—  2)  was  taken  as  the  standard  for  molecular  masses. 

257.  Laws   of  Simple  and  Multiple   Volumes. - 

Craseous  bodies  combine  in  definite  proportions  not  only 
by  weight,  as  do  liquids  and  solids  as  well,  but  also  by 
volume.  This  was  the  case  when  oxygen  and  hydrogen 
united  to  form  steam  (cf.  §  38),  and  when  hydrogen  and 
chlorine  gave  hydrochloric  acid  gas  (cf.  §  94).  The  re- 
lation between  the  volumes  of  the  combining  gases  and 
the  volumes  of  the  products,  if  these  are  gaseous,  is  also 
simple  and  definite.  Thus,  two  volumes  of  hydrogen 
and  one  of  oxygen  give,  on  union,  two  volumes  of  steam  ; 
and  three  volumes  of  hydrogen  and  one  of  nitrogen 
give  two  of  ammonia. 

These  and  other  facts  give  a  basis  for  the  Law  of 
Simple  Volumes,  which  is  :  The  volumes  of  reacting  gase- 
ous substances  have  a  simple  relation  with  one  another 
and  with  the  volumes  of  the  products,  if  these  are  gaseous. 

Just  as  we  have  a  Law  of  Multiple  Proportions  by 
Weight,  so  we  have  a  Law  of  Multiple  Volumes.  It 
is  stated  thus :  If  two  gases  unite  in  more  than  one  pro- 
portion, and  we  consider  the  volume  of  one  as  fixed,  then 
the  several  volumes  of  the  second  gas  that  unite  with  the 
fixed  volume  of  the  first  are  in  a  simple  relation  to  one 
another. 

Thus,  the  volumes  of  nitrogen  which  unite  with  one  volume 
of  oxygen  to  form  nitrous  oxide  (N2O)  and  nitric  oxide  (NO) 
are  as  2  : 1  respectively  (cf.  §§  170  and  171). 


246  MOLECULES  AND  ATOMS. 

258.  Volumetric  Meaning  of  an  Equation.  —  Since 
the  formula  of  a  gaseous  substance  means  a  molecule  of 
the  substance,  it  means  also  a  volume  of  it;  for  equal 
volumes  contain  approximately  equal  numbers  of  mole- 
cules. 

Thus,  in  the  equation, 

2  II2  +  02  =  2  H20, 

H2  means  one  volume  of  hydrogen;  and  2  H2,  two  vol- 
umes. O2  means  one  volume  of  oxygen;  2  H9O  means 
two  volumes  of  steam. 

So  the  equation  N2  -f-  3  H2  =  2  NH3  means  that  one 
volume  of  nitrogen  and  three  volumes  of  hydrogen  unite 
to  give  two  volumes  of  ammonia.  These  relations  have 
all  been  proved  by  experiment. 

From  a  correctly  written  equation  we  can,  therefore, 
know  in  what  proportions  by  volume  gaseous  substances 
unite. 

Thus,  the  equation, 

2  CO  +  O2  =  2  CO2, 

shows  that  one  volume  of  oxygen  unites  with  two  of  carbon 
monoxide,  forming  two  of  carbon  dioxide.  In  other  words, 
the  volume  of  the  carbon  dioxide  formed  by  burning  a  given 
volume  of  carbon  monoxide  is  just  equal  to  the  volume  of  the 
carbon  monoxide. 
The  equation, 

CH4  +  2  O2 »  CO2  -f  2  H2O, 

(c/.  §  218)  shows  that  one  volume  of   marsh  gas  uses  up  two 


VALENCE.  247 

volumes  of  oxygen  in  forming  one  volume  of  carbon  dioxide 
and  two  volumes  of  steam. 

If  a  substance  that  appears  in  a  reaction  is  not 
known  in  the  gaseous  state,  nothing  can  be  said  of  its 
volume. 

Thus,  the  equation,  C  -\-  O2 >  CO2,  tells  us  that  one  vol- 
ume of  oxygen  disappears  in  the  formation  of  one  volume  of 
carbon  dioxide  ;  but  it  does  not  tell  us  the  volume  of  the  car- 
bon that  unites  with  one  volume  of  oxygen,  since  we  cannot 
experiment  with  gaseous  carbon. 

259.  Valence.  —  The  student  must  have  noticed 
from  the  formulas  previously  studied  that  atoms  differ 
greatly  in  their  power  of  combining  with  other  atoms. 
Thus,  the  formulas  of  the  compounds  of  hydrogen  show 
interesting  differences ;  for,  while  in  the  case  of  hydro- 
chloric acid  one  atom  of  chlorine  unites  with  one  of 
hydrogen,  in  the  case  of  water  (H2O)  one  atom  of 
oxygen  holds  two  of  hydrogen.  The  combining  powers 
of  the  nitrogen  and  the  carbon  atom  are  still  greater ; 
for  in  the  molecule  of  ammonia  (NH8)  one  nitrogen 
atom  holds  three  hydrogen  atoms ;  and  in  the  molecule 
of  marsh  gas  (CH4)  one  carbon  atom  holds  four  atoms 
of  hydrogen. 

This  power  of  the  atoms  to  unite  with  different  num- 
bers of  other  atoms  is  called  valence.  An  element  like 
chlorine,  whose  atom  can  hold  only  one  atom  of  hydro- 
gen, is  said  to  have  a  valence  of  one,  or  to  be  univa- 


248  MOLECULES  A.VD  ATOMS. 

lent.      The  hydrogen  atom  is  always  considered  univa- 

lent. 

• 

An  element  like  oxygen,  whose  atom  can  hold  two  hydrogen 
atoms,  is  said  to  have  a  valence  of  two,  or  to  be  bivalent.  One 
like  nitrogen  is,  therefore,  trivalent  J  and  one  like  carbon, 
quadrivalent. 

Valence  is  not  only  the  power  of  combining  with,  but 

also  of    replacing,  different  numbers  of    atoms.     Thus, 

i£ 
the  formula  of  potassium  sulphate,     ^  SO4,  is    derived 

TT 

from  that    of    sulphuric    acid,   ,  j  SO4,  by  replacing  two 

hydrogen  atoms  by  two  of  potassium.  Potassium  has, 
therefore,  a  valence  of  one. 

In  the  case  of  zinc  nitrate,  Zn  XTr»3'  one  atom  of  zinc  has  re- 

JMJg, 

placed  two  hydrogen  atoms  (in  2  HNO8)  ;  hence  zinc  is  bivalent. 
Similarly,  iron  is  trivalent  in  ferric  phosphate,  FePO4 ;  for  the 
formula  of  phosphoric  acid  is  H8PO4.  Aluminum,  also,  is 
trivalent  in  aluminum  sulphate,  A12  (SO4)3  ;  for  two  aluminum 
atoms,  each  with  a  valence  of  three,  replace  six  hydrogen 
atoms  (in  3  H2SO4). 

The  valence  of  an  element  is  often  represented  by 
small  Roman  figures  placed  a  little  above  and  to  the 
right  of  the  symbol  of  the  element.  Thus,  A  I111  means 
trivalent  aluminum ;  Hg1,  univalent  mercury ;  and  PtIV, 
quadrivalent  platinum. 

260.  Different  Formula  Types  Based  on  Valence. 
—  A  bivalent  element,  like  oxygen,  unites  with  two 


GRAPHIC  FORMULAS.  249 

atoms  of  a  univalent  element,  but  with  only  one  of  an- 
other bivalent  element.  Thus,  calcium  chloride  has  the 
formula  CaCl2 ;  but  calcium  oxide  has  the  formula  CaO. 

When  a  bivalent  element  unites  with  a  trivalent  ele- 
ment, two  atoms  of  the  trivalent  element  generally 
require  three  of  the  bivalent  one.  This  is  shown  in 
the  formulas  A12O3  for  aluminum  oxide,  As2S3  for 
arsenious  sulphide,  and  Mg8N2  for  magnesium  nitride. 

When  quadrivalent  atoms,  like  those  of  carbon  and 
silicon,  unite  with  univalent  atoms,  four  of  the  univa- 
lent atoms  are  required,  as  in  CH4  and  SiCl4;  when 
they  unite  with  bivalent  atoms,  two  of  the  latter  are 
usually  required,  as  in  CS2  and  SiO9. 

The  valence  of  an  element  is  not,  however,  fixed  ;  for  car- 
bon forms  the  compound  CO,  in  which  its  valence  is  undoubt- 
edly two.  So,  also,  nitrogen  is  trivalent  in  the  compounds 
NTI3  and  N2O3,  but  quinquivalent  in  NH4C1  and  N2O5. 

261.  Graphic  Formulas. — By  means  of  the  idea  of 
valence,  we  may  represent  the  relation  of  atoms  to  one 
another  in  molecules.  When  the  molecule  consists  of  two 
atoms,  only  one  arrangement  is  possible,  viz.,  the  atoms 
are  joined  directly.  Thus  in  hydrochloric  acid  and  in 
molecular  hydrogen  we  have  simple  union.  If  we  repre- 
sent the  combining  power  of  the  elements  by  lines 
(called  bonds),  we  may  write  the  formula  for  hydro- 
chloric acid  graphically  H —  Cl  and  that  of  hydrogen 
H  —  H.  The  single  line  shows  that  the  valence  of  each 
atom  in  the  molecule  is  one, 


250  MOLECULES  AND   ATOMS. 

Everything  goes  to  show  that  the  two  hydrogen  atoms  in 
the  water  molecule  are  not  united  to  each  other,  but  to  oxygen. 
Similarly  it  is  believed  that  the  hydrogen  atoms  of  the  am- 
monia molecule  are  all  united  to  nitrogen,  and  those  of  the 
marsh-gas  molecule  to  carbon.  We  may  represent  these  facts 


in  the  formulas,  — 


H 


^11  | 

II  —  O—  II  ;  ;  and  II  —  C  —  II. 


Formulas  like  those  just  given  are  called  graphic  01 
structural  formulas.  We  may  represent  an  element  in 
its  nascent  state  by  the  symbol  with  free  valence  bonds. 
Thus,  —  O  —  represents  nascent  or  atomic  oxygen  ; 
H  — ,  nascent  hydrogen. 

262.  Isomerism. — Graphic  formulas  enable  us  to 
represent  differences  between  compounds  which  cannot 
be  distinguished  by  the  ordinary  formulas.  Methyl 
ether  and  ethyl  alcohol,  for  example,  are  both  repre- 
sented by  the  formula  C9H6O ;  these  substances  are  so 
different,  however,  that  no  one  would  mistake  one  for 
the  other.  Thus,  methyl  ether  boils  at  — 24°  C.,  at  or- 
dinary pressure,  while  ethyl  alcohol  boils  at  -)-78°  C. 
Such  compounds  are  said  to  be  isomeric  with  each 
other. 

By  the  use  of  graphic  formulas,  the  difference  in  character 
between  these  isomeric  substances  is  readily  understood. 
Thus,  the  graphic  formula  for  methyl  ether  is, — 


ALLOTROPISM.  251 

H  II 

I  I 

H— C— O— C— H; 

I  I 

II  H 

while  that  of  alcohol  is,  — 

II      II 

I        I 
H  — C  — C  — O— H. 

I        I 
II     H 

According  to  these  formulas,  all  the  hydrogen  atoms  of 
methyl  ether  have  the  same  relation  to  the  remainder  of  the 
molecule  and  should  behave  in  the  same  way  with  reagents; 
while  in  the  case  of  ethyl  alcohol  one  hydrogen  atom  —  the 
one  bound  to  oxygen  —  should  be  different  from  the  other 
five.  This  is  actually  the  case;  for  the  atom  of  hydrogen 
bound  to  oxygen  is  the  only  one  of  the  six  that  can  be  re- 
placed by  sodium  and  other  metals. 

263.  Allotropism.  —  Just  as  there  are  compounds 
having  the  same  chemical  composition  which  are  yet  very 
unlike  in  their  properties,  so  there  are  elements  existing 
in  forms  so  different  that  they  might  easily  be  sup- 
posed to  be  entirely  different  substances.  In  the  pre- 
ceding section  we  have  called  the  compounds  isomers ; 
the  different  forms  of  the  same  element  are  called  allo- 
tropic  forms  of  the  element.  The  existence  of  an 
element  in  different  forms  is  called  allotropism.  Car- 
bon and  sulphur,  as  we  have  already  learned,  exist  in 
several  allotropic  forms  ;  the  same  is  true  of  oxygen, 
phosphorus,  silicon,  boron,  etc. 


252  MOLECULES  AND  ATOMS. 

Allotropism  is  probably  due  to  different  causes,  such  as 
different  arrangements  of  the  atoms  in  the  molecule,  or  differ- 
ent numbers  of  atoms  in  the  molecule. 

In  many  cases  an  allotropic  form  is  only  temporary.  This 
is  true  of  the  plastic,  or  amorphous,  modification  of  sulphur 
(cf.  §  175),  which  changes  into  the  ordinary  form  with  evolution 
of  heat.  Plastic  sulphur  thus  represents  a  condition  of  unstable 
equilibrium,  like  a  compound  which  has  a  negative  heat  of  for- 
mation (cf.  §  235). 

264.  Exercises. 

1.  (a)  A  liter  of  a  certain  gaseous  substance  weighs  approxi- 
mately 1.966  grams  at  standard  conditions  ;  what  is  its  molecu- 
lar mass  ?  (See  §  242.) 

(6)  If  -ft-  of  this  substance  is  carbon  and  -ft  oxygen,  what 
is  its  formula  ?  (See  §  252.) 

2.  (a)  Two  hundred  c.c.  of  a  gas  weigh  0.3932   grams  at 
standard  conditions  ;  what  is  the  molecular  mass  ? 

(6)  Analysis  shows  that  the  gas  is  composed  of  nitrogen, 
63.64%,  and  oxygen,  36.36% ;  what  is  the  formula? 

3.  The   oxide   of    magnesium  is  composed  of  magnesium, 
60%,  oxygen,  40%;  what  is  its  simplest  formula? 

4.  A   chloride    of   phosphorus   has   the   composition,  phos- 
phorus 22.55%,  chlorine  77.45%  ;  find  its  simplest  formula. 

5.  What  is  the  approximate  atomic  mass  of  platinum  if  its 
specific  heat  is  about  0.033  ? 

6.  What  volume   of   oxygen   is   used v  up   when  20   c.c.  of 
acetylene  burn  in  air  ?     What  is  the  volume  of  carbon  dioxide 
formed  ?     (Cf.  §§  221  and  258.) 

7.  Write  the  molecular  equation  for  the  combustion  of  pen- 
tane,  C5H10,  in  oxygen,  if  the  products  are  carbon  dioxide  and 
water. 

What  volume  of  oxygen  is  used  up  when  50  c.c.  of  pentane 
burn? 


EXERCISES.  253 

What  volume  of  carbon  dioxide  is  produced  ? 

8.  Knowing  that  the  valence  of  an  element  x  is  1,  write 
the  simplest  formula  for  its  sulphate,  its  carbonate,  and  its 
nitrate. 

Q.  Write  the  simplest  formulas  of  the  chloride  and  sulphite 
of  an  aliment  whose  valence  symbol  is  Siri. 


CHAPTER   XVII. 
FLUORINE,  BROMINE,  IODINE,  AND  THEIR  COMPOUNDS. 

265.  Halogens.  —  The    elements    fluorine,    chlorine, 
bromine,    and    iodine    are    called  "  the   halogens,"  from 
hah,    Greek   for  "salt,"  and    the    suffix  gen,    meaning 
"  a  constituent  of,"  as  in  "  hydrogen,"  etc. 

Fluorine  and  chlorine  are  gaseous  at  the  ordinary  tempera- 
ture ;  bromine  is  a  liquid  boiling  at  about  50°  C.  ;  while  iodine 
•is  a  black  solid  which  gives  off,  even  at  ordinary  temperatures, 
a  beautiful,  violet  vapor. 

266.  Fluorine.  —  The   element  fluorine  was   known 
in  its  compounds  long  before  it  was  obtained  in  the  free 
condition.     The  most  common  of  its  compounds  is  cal- 
cium fluoride,  or  fluorspar,  CaFl2.     Fluorspar  derives  its 
name  frqmfliw,  Latin  for  "to  flow,"  and  spar,  meaning 
"  a  rock."     The  name  is  applied  to  this  substance  owing 
to  the  use  of  fluorspar  as  'A  flux  in  metallurgy. 

A  flux  is  an  easily  fusible  substance  added  to  the  mixture 
of  an  ore  and  a  reducing  agent  to  promote  fusion  of  the  mix- 
ture. The  substance  resulting  from  the  union  of  the  flux  with 
the  impurities  present  is  usually  called  "  slag." 

Another  important  natural  fluorine  compound  is  cryo- 
lite. This  is  a  double  fluoride  of  aluminum  and  sodium; 
its  formula  is  A1FL.  3  NaFl  or  NaQAlFL. 

o  o  o 

254 


HYDROFLVOHIC  ACID.  255 

Fluorine  cannot  be    prepared    from    either    of  these 
compounds  directly,  but  has  been  made  (1886)  by  the 
electrolysis  of  anhydrous  hydrofluoric  acid,  HFL 
2  HF1  =  H2  +  F12. 

The  operation  may  be  carried  out  in  copper  or  plati- 
num apparatus,  but  not  in  glass,  since  hydrofluoric  acid 
attacks  glass  energetically. 

267.  Properties  of  Fluorine.  —  Fluorine  is  a  yellow 
gas,  about  one  and  two-fifths  times  as  heavy  as  air.     It 
acts  upon  water  with  violence,  according  to  the  equation, 

2  H2O  +  2  Fig »  4  HF1  +  O2. 

The  oxygen  formed  always  contains  some  ozone  (c/.  §  287). 

Fluorine  unites  with  hydrogen  explosively,  even  in 
the  dark  (cf.  Chlorine,  §  84),  to  form  hydrofluoric  acid. 
It  forms  no  compounds  with  oxygen,  so  far  as  known. 
As  commonly  prepared,  fluorine  acts  upon  glass,  but 
this  is  due  to  the  fact  that  a  small  amount  of  hydro- 
fluoric acid  is  present  in  the  fluorine. 

Fluorine  acts  readily  upon  silicon  and  antimony, 
forming  the  corresponding  fluorides,  SiFl4  and  SbFl3. 

Gaseous  fluorine  has  been  condensed  to  a  liquid  boiling  at 
— 187°  C.  at  ordinary  pressure. 

268.  Hydrofluoric  Acid.  — Hydrofluoric  acid  is  com- 
monly prepared  by  heating  calcium  fluoride  with  con- 
centrated sulphuric  acid.     The  equation  is,  - 

CaFl2  +  H2S04  =  CaSO4  +  2  HFL 


256  FLUORINE,   BROMINE,   IODINE. 

Anhydrous  hydrofluoric  acid  is  a  liquid  boiling  at 
about  19°  C. 

Both  the  vapor  of  this  liquid  and  its  solution  in  water  are 
very  poisonous.  The  aqueous  solution  reacts  with  almost  all  the 
metals,  forming  fluorides  and  hydrogen  ;  and  decomposes  the 
oxides,  forming  fluorides  and  water. 

Silicon  dioxide  (quartz,  sand,  etc.)  gives  with  hydro- 
fluoric acid  silicon  tetrafluoride  (SiFl4)  and  water,  ac- 
cording to  the  equation, 

SiO2  +  4  HF1  =  SiFl4  +  2  H2O. 

Silicon  tetrafluoride  is  a  gas. 

Glass  —  a  mixture  of  silicates,  i.e.,  salts  of  silicic 
acid,  H2SiO3  —  is  acted  upon  by  hydrofluoric  acid  as 
silicon  dioxide  is.  Thus,  we  may  represent  the  action 
of  calcium  silicate,  CaSiO3,  with  the  acid  by  the  equa- 
tion, 

CaSiO3  +  6  HF1  =  CaFl2  +  SiFl4  +  3  H2O. 

Hence,  when  glass  is  treated  with  hydrofluoric  acid,  the  sili- 
con present  in  the  glass  escapes  as  SiFl4,  leaving  a  depression 
in  the  glass.  This  fact  is  made  use  of  in  the  operation  of  etch- 
ing glass.  The  glass  is  first  covered  with  a  thin  layer  of  paraf- 
fin, and  a  design  is  drawn  in  the  paraftin  by  means  of  a  sharp 
point.  When  the  exposed  glass  is  wet  with  the  solution  of  the 
acid  (a  swab  of  cotton  attached  to  a  stick  may  be  used  to  ap- 
ply the  solution),  or  is  left  in  the  vapor  of  the  acid,  the  design 
is  etched  into  the  glass. 

Hydrofluoric  acid  is  commonly  kept  in  bottles  of 
paper,  covered  inside  and  out  with  a  thick  layer  of  par- 


PREPARATION   OF  BROMINE.  257 

affin.     Vessels  of  lead,  platinum,  or  rubber  may  also  be 
used. 

Bromine. 

269.  Preparation  of  Bromine.  —  Bromine  is  found 
in  nature  in  the  combined  form,  chiefly  as  bromides.  The 
most  common  bromides  are  those  of  sodium  (NaBr),  of 
potassium  (KBr),  and  of  magnesium  (MgBr2). 

Bromides  occur  in  sea-water  and  in  connection  with 
salt  deposits. 

Bromine  is  prepared  hy  heating  a  bromide  with  man- 
ganese dioxide  and  dilute  sulphuric  acid.  The  bromine 
vapor  evolved  is  condensed  in  cold  receivers.  With 
sodium  bromide  the  equation  is,  — 

Mn02-|-  2  NaBr  -f  3  H2SO4  =  MnSO4  +  2  NaHSO4  -f  Br2  -f 
2H20. 

This  reaction  is  like  that  used  in  making  chlorine  (cf.  §81) 
from  common  salt,  manganese  dioxide,  and  sulphuric  acid. 
The  reaction  takes  place  in  at  least  two  stages  :  — 

(1)  The  sulphuric  acid  and  sodium  bromide  give  sodium  hy- 
drogen sulphate  and  hydrobromic  acid,  according  to  the  equa- 
tion, 

NaBr  +  H2S04 »  NaHSO4  +  HBr  ;  and 

(2)  The  hydrobromic  acid  and  the  manganese  dioxide  react 
to  give  manganous  bromide  (MnBr2). 

MnO2  +  4  HBr  =  MnBr2  +  Br2  -f  2  H2O. 

The  manganous  bromide  and  sulphuric  acid  then  give  rise 
to  manganous  sulphate  (see  above)  and  more  hydrobromic  acid. 

MnBr2  +  H2SO4 >  MnSO4  +  2  HBr. 


258  FLUORIXE,   BROMINE,   IODINE. 

Bromine  may  also  be  prepared  by  conducting  the 
proper  amount  of  chlorine  into  the  solution  of  a  brom- 
ide. With  magnesium  bromide  the  equation  is,  — 

MgBr,  +  C13 >  Mg('l2  +  Brr 

270.  Properties  of  Bromine.  —  Bromine  is  a  brown 
liquid  about  3.2  times  as  heavy  as  water.  Its  vapor  has 
an  odor  much  like  that  of  chlorine,  and  affects  the  eyes. 
Bromine  boils  at  about  59°  C. 

The  density  of  bromine  vapor  shows  that  the  mole- 
cule is  diatomic  ;  its  formula  is,  therefore,  Br2.  At  about 
1000°  C.  the  molecule  begins  to  dissociate  into  molecules 
containing  only  one  atom  each  (cf.  §  45). 

Bromine  dissolves  in  water,  carbon  disulphide,  and 
other  solvents.  The  aqueous  solution  is  called  "  brom- 
ine water." 

In  the  presence  of  some  substance  capable  of  taking 
up  oxygen,  bromine  reacts  with  water  energetically, 
according  to  the  equation, 

H20  +  Br2 >  2  HBr  +  (—  O  -). 

By  —  O  —  we  mean  nascent  oxygen  (cf.  §  261). 
Bromine  water  is  thus  a  good  oxidizing  agent. 

The  same  action  goes  on  more  slowly  when  no  oxidizable 
substance  is  present ;  bromine  water  thus  becomes  converted 
into  a  dilute  solution  of  hydrobromic  acid,  HBr. 

Bromine  is  less  active  than  chlorine,  but,  like  chlor- 
ine, it  unites  with  hydrogen  and  with  metals  to  form 
bromides. 


PROPERTIES   Of    HYDROBROMIC  ACID.  259 

271.  Hydrobromic  Acid.  —  Hydrobromic  acid  cannot 
be  made  in  a  pure  state  by  treating  a  bromide  with  con- 
centrated sulphuric  acid,  for  the  reason  that  some  of  the 
hydrobromic  acid  formed  breaks  up  into  hydrogen  and 
bromine,  and  the  nascent  hydrogen  reduces  the  sulphuric 
acid.     The  products  of  the  action  are  thus  bromine  and 
sulphurous  acid,  as  well  as  hydrobromic   acid.     These 
facts  are  represented  in  the  equations, 

(1)  NaBr  4-  H2SO4 >  NaHSO4  -f  HBr  ; 

(2)  2  HBr  +  H2S04 >  H2O  +  Br2  +  H2SO3. 

The  method  commonly  used  to  prepare  hydrobromic 
acid  is  to  treat  red  phosphorus  with  bromine  in  the 
presence  of  water.  The  phosphorus  and  bromine  first 
unite  to  form  phosphorus  tribromide,  PBr3;  but  this 
substance  is  decomposed  at  once  by  the  water  to  form 
phosphorous  acid  and  hydrobromic  acid.  The  equations 

are,— 

(1)  2P+3Br2:=2PBr3; 

(2)  PBr3  +  3  H20  =  H3PO3  +  3  HBr. 

The  phosphorous  acid,  being  non-volatile,  remains  behind ; 
while  the  gaseous  hydrobromic  acid  passes  off.  The 
hydrobromic  acid  may  be  freed  from  bromine  vapor 
by  passing  it  through  a  U-tube  containing  moist  red 
phosphorus. 

272.  Properties    of    Hydrobromic    Acid. —  Hydro- 
bromic  acid  gas  is  like  hydrochloric  acid  gas.     It  fumes 
in  the  air  and  dissolves  readily  in  water.     Its  concen- 


260  FLUORINE,   BROMINE,   IODINE. 

trated  aqueous  solution  has  a  specific  gravity  of  almost 
1.8,  and  contains  82J&  by  weight  of  the  acid. 

Hydrobromic  acid  begins  to  dissociate  into  its  elements  (c/. 
§  45)  at  about  800°  C.;  it  is,  therefore,  much  less  stable  than 
hydrochloric  acid,  which  begins  to  dissociate  at  about  1500°  C. 

Iodine. 

273.  Occurrence  and  Preparation  of  Iodine.  —  The 

chief  source  of  iodine  until  recently  was  the  ashes  of 
certain  sea-plants  which  absorb  iodine  compounds  from 
sea-water.  At  the  present  time  the  element  is  obtained 
largely  from  the  Chile  saltpeter  deposits.  In  these 
deposits  the  iodine  is  found  chiefly  as  sodium  iodate, 
NaI03. 

Iodine  may  be  set  free  from  an  iodide  in  just  the  same 
way  that  chlorine  and  bromine  are  set  free  from  chlorides 
and  bromides  respectively,  viz.,  by  heating  it  with  a 
mixture  of  manganese  dioxide  and  dilute  sulphuric 
acid.  A  representative  equation  is,  — 

MnO2  +  2  Nal  +  3  H2SO4  =  MnSO4  +  2  NaHSO4  + 
2H20  +  I2. 

The  stages  in  which  the  reaction  takes  place  are  partly 
represented  by  the  equations, 

(1)  Kal  +  H2SO4  =  HI  -f  NaHSO4; 

(2)  Mn02  +  4  HI  =  MnI2  +  2  H2O  +  2  I; 

(3)  MnI2  +  H2S04  =  MnS04  +  2  HI. 

Iodine  may  also  be  set  free  from  an  iodide  by  means 
of  chlorine  or  bromine  (cf.  §  269). 


HYDRIODIC  ACID.  261 

2  Nal  +  Br2 »  2  NaBr  +21. 

2  Nal  +  C12" »  2  NaCl  +  2  I. 

274.  Properties  of  Iodine.  —  Iodine  is,  ordinarily,  an 
almost  black  solid,  melting  at  114°  C.  and  boiling  at 
about  184°  C.     Its  vapor  has  a  beautiful,  violet  color; 
it  is  about  8.7  times  as  heavy  as  air. 

Iodine  is  very  soluble  in  carbon  disulphide  and  in 
ether,  but  only  slightly  soluble  in  water.  It  is  less 
active  than  chlorine  or  bromine.  It  stains  the  skin 
brown,  and  imparts  an  intensely  blue  color  to  starch 
paste.  Iodine  sublimes  (ef.  §  149)  when  heated.  » 

Up  to  about  600°  C.  the  molecule  of  iodine  vapor  consists 
of  two  atoms ;  above  this  temperature,  dissociation  takes 
place.  At  about  1500°  C.  only  monatomic  iodine  molecules 
exist. 

275.  Hydriodic  Acid.  —  Hydriodic  acid  is  still  more 
unstable  than  hydrobromic  acid.     It  cannot  be  made  in 
a  pure  condition  by  treating  an  iodide  with  concentrated . 
sulphuric  acid  for  the  reason  that  the  hydriodic  acid 
reduces   the    sulphuric    acid.     The   reduction  goes  not 
only  to  sulphurous  acid,  as  in  the  case  of  hydrobromic 
acid,  but  in  part  even  to  hydrogen  sulphide.     The  equa- 
tions representing  this  are,  — 

(1)  KI  +  H2SO4  =  KHSO4  +  HI ; 

(2)  2  HI  +  H2S04  =  H2S03  +  H0O  +  21; 
(2a)  8  HI  +  H2S04  =  H2S  +  4  H2O  +  81. 

Hydriodic  acid  gas  may  be  made  by  allowing  red 
phosphorus  and  iodine  to  react  in  the  presence  of  water. 


FLUORINE,   BROMINE,   IODINE. 

Phosphorus  tri-iodide  is  first  formed,  but  is  decomposed 
at  once  according  to  the  equation, 

PIS  +  3  H20  =  H8P08  +  3  HI  (cf.  §  271). 

An  aqueous  solution  of  hydriodic  acid  may  best  be 
prepared  by  making  use  of  a  property  common  to 
chlorine,  bromine,  and  iodine,  viz.,  the  ability  of  each 
of  these  substances  to  decompose  hydrogen  sulphide. 
The  sulphur  formed,  being  insoluble,  is  precipitated ; 
hence  the  reaction  goes  on  to  completion  (cf.  §  180). 
The  equation  in  the  case  of  iodine  is, 

2  I  -f  H2S >  2  HI  +  S. 

The  hydrogen  sulphide  is  conducted  into  a  mixture  of 
iodine  and  water  until  the  iodine  disappears.  The  sulphur  is 
then  filtered  off,  and  the  filtrate  distilled.  After  the  water  has 
passed  off,  a  heavy  liquid  is  obtained,  which  boils  at  126°  0. 
This  is  about  57%  hydriodic  acid. 

276.  Properties  of  Hydriodic  Acid.  —  Hydriodic 
acid  gas  is  about  4.4  times  as  heavy  as  air.  Like  hy- 
drochloric and  hydrobromic  acids,  it  is  very  soluble  in 
water.  One  cubic  centimeter  of  water  at  10°  C.  and 
standard  pressure  dissolves  about  450  c.c.  of  the  gas. 

Hydrogen  and  iodine  can  be  made  to  unite  under 
appropriate  conditions.  When  uniting  they  do  not 
evolve  heat,  but  absorb  it.  This  accounts  for  the  fact 
that  hydriodic  acid  is  so  unstable  (cf.  §  235). 

The  dissociation  of  hydriodic  acid  is  like  that  of 
steam  (cf.  §  45).  At  any  temperature  above  the  point 


COMPOUNDS  OF  THE  HALOGENS  WITH  OXYGEN.     263 

at  which  dissociation  begins,  the  decomposition  of  hydri- 
odic  acid  into  hydrogen  and  iodine  goes  on  side  by  side 
with  recombination  of  hydrogen  and  iodine  to  form  hy- 
driodic  acid.  The  condition  of  equilibrium  is  reached 
when  as  many  molecules  of  hydriodic  acid  are  formed 
in  a  given  time  as  are  decomposed  in  the  same  time. 
We  may  represent  this  condition  of  equilibrium  by  the 
equation, 

2  HI  jn±  H2  +  I2. 

Such  an  equation  is  called  a  "balanced"  or  "equilib- 
rium "  equation.  The  arrows  indicate  that  the  reac- 
tion goes  in  both  directions. 

If  either  product  of  dissociation  is  removed  from  the 
"  sphere  of  action,"  the  dissociation  goes  on  to  comple- 
tion. Thus,  if  silver  is  placed  in  hydriodic  acid  solu- 
tion, it  unites  with  the  iodine  as  rapidly  as  iodine  is 
formed.  Hence  hydrogen  is  set  free. 

2  Ag  +  2  HI »  2  Agl  +  H2. 

Because  of  its  ready  dissociation,  hydriodic  acid  acts  as  a 
powerful  reducing  agent.  Oxygen,  or  any  oxidizing  agent,  gives 
with  it  iodine  and  water. 

4  HI  +  02 »  2  H20  +  4  I. 

277.  Compounds  of  the  Halogens  with  Oxygen.  — 

Only  three  halogen  oxides  have  been  actually  made ; 
although  more  —  especially  one  oxide  of  bromine  —  are 
suspected  to  be  capable  of  existing.  The  three  oxides 
definitely  known  are :  — 


264  FLUORINE,   BE 0 MINE,   IODINE. 

Chlorine  monoxide,  C12O  ; 
Chlorine  dioxide,  C1O2  or  C12O4; 
Iodine  pentoxide,  I2O&. 

Chlorine  monoxide  is  an  unstable  liquid  which  boils 
at  -)-5°  C.  It  explodes  with  violence.  When  one 
volume  of  chlorine  monoxide  is  carefully  decomposed,  it 
gives  half  a  volume  of  oxygen  and  one  volume  of 
chlorine. 

Chlorine  dioxid^  (C1O2)  and  tetroxide  (C12O4)  corre- 
spond to  the  nitrogen  oxides  NO2  and  N2O4.  It  is 
formed  with  violent  explosion  when  concentrated  sul- 
phuric acid  acts  upon  potassium  chlorate.  The  prepara- 
tion of  this  substance  should  not  be  attempted  without 
precise  directions  and  extraordinary  precautions. 

Chlorine  dioxide  is  a  reddish-brown  liquid,  boiling  at 
about  10°  C. 

Iodine  pentoxide  is  the  only  oxide  of  iodine  known  at 
present.  It  is  a  white,  stable  powder;  with  water  it 
gives  iodic  acid, 

I206  +  H20  =  2  HI03. 
Iodine  pentoxide  is  thus  the  anhydride  of  iodic  acid. 

278.  Compounds  of  the  Halogens  with  Oxygen  and 
Hydrogen.  — Compounds  of  the  halogens  with  oxygen 
and  hydrogen  are  called  oxygen  acids,  or  oxy-acids,  of 
the  halogens.  They  are  more  numerous  than  com- 
pounds with  oxygen  alone ;  for  they  include  at  least 
three  chlorine  acids,  three  bromine  acids,  and  two  iodine 


HYPOCHLOROUS  ACID.  265 

acids.  The  formulas  of  these  acids  appear  in  the  fol- 
lowing table :  — 

HC10.  IIBrO. 

IiriO    (  known  only  | 
1U2  I  in  its  salts     J 

HC103.  HBr03.  HIO3. 

HC1O4.  HBrO4.  HIO4. 

279.  Hypochlorous  Acid,  H-0-C1.  —  Hypochlorous 
acid  is  present  in  a  solution  of  chlorine  in  water  (<?f. 
§85). 

The  equation  for  the  action  of  chlorine  upon  water  is  a 
"balanced"  one,  viz., 

H20  +  C12  ~ »  HOC1  +  HC1 ; 

there  is,  therefore,  an  equilibrium  between  these  four  substances. 
If  a  substance  is  present  which  is  capable  of  taking  up  oxygen, 
the  products  of  the  action  are  hydrochloric  acid  and  nascent 
oxygen,— 

HOC1  +  HC1 >  2  HC1  +  (—  O  — ). 

In  the  presence  of  sunlight  a  concentrated  solution  of  chlor- 
ine in  water  gives  off  oxygen  (c/.  §  85). 

Salts  of  hypochlorous  acid,  i.  e.,  hypochlorites,  are 
formed  by  passing  chlorine  into  dilute  solutions  of 
hydroxides  of  the  metals.  The  amount  of  chlorine 
used  must  be  less  than  is  required  to  saturate  the  hy- 
droxide. With  potassium  hydroxide  the  equation  is, — 

2  KOH  +  C12 »  KOC1  +  KC1  +  H2O. 


266  FLUORINE,    BROMINE,   IODINE. 

When    chlorine     acts    upon    powdered    slaked    lime, 
Ca(OH)2,  Ueaching  powder  (Ca^p,     \    is  produced. 

C1  /OC1 


280.  Chlorous  Acid,  H-0-C1=0.  —  Chlorous  acid  does 
not  exist  free.  Its  potassium  salt  is  produced,  iu  solution,  when 
an  aqueous   solution  of  chlorine  dioxide,  C1O2,  is  treated  with 
potassium  hydroxide. 

281.  Chloric     Acid,    H-O-Cl^Q.  —  Chloric    acid    is 

known  only  in  its  aqueous  solution  ;  this  may  be  con- 
centrated until  it  contains  40^  of  chloric  acid.  Chloric 
acid  is  a  powerful  oxidizing  agent. 

Chlorates  are  formed  by  conducting  chlorine  into  hot, 
concentrated  solutions  of  alkalies  to  complete  saturation. 
With  potassium  hydroxide  the  equation  is,  — 

6  KOH  +  3  C12  -  »  KC1O3  +  5  KC1  +  3  H2O. 

The  chlorate  is  separated  from  the  chloride  by  recrystalliza- 
tion  from  hot  water.  The  potassium  chlorate,  being  much  less 
soluble  than  the  other,  separates  out  first. 

Potassium  chlorate  has  already  been  used  to  produce 
oxygen  (§  19).  The  equations  representing  the  action 
of  hydrochloric  acid  upon  potassium  chlorate  (§82) 
are,  in  part,  — 

(1)  KC1O3  +  HC1  -  »  KC1  +  HC1O3; 

(2)  HC108  -j-  5  HC1  -  >  3  H20  +  3  Cla. 


COMPOUNDS  OF  BROMINE  WITH  OXYGEN,   ETC.    267 

^° 

282.  Perchloric  Acid,  H-0-C1=0. —  Perchloric  acid 

^0 

is  a  colorless,  explosive  liquid  about  1.8  times  as  heavy 
as  water.  Its  salts,  the  perchlorates,  are  produced  when 
the  chlorates  are  partly  decomposed  by  heat. 

In  the  decomposition  of  potassium  chlorate  (c/.  §  19),  a  point 
is  soon  reached  at  which  a  considerable  increase  of  temperature 
is  needed  to  continue  the  evolution  of  oxygen.  The  amount  of 
oxygen  evolved  up  to  this  point  is  only  one-third  of  the  quantity 
present  in  potassium  chlorate.  If  we  stop  at  this  stage,  we 
obtain  a  mixture  of  potassium  perchlorate  and  potassium  chlor- 
ide, as  shown  in  the  equation, 

2  KC1O3  =  KC1O4  +  KC1  +  O2. 

The  potassium  chloride  is  much  more  soluble  than  the  per- 
chlorate; hence  these  substances  may  be  separated  by  recrystal- 
lization  from  water. 

283.  Compounds  of  Bromine  with  Oxygen  and  Hy- 
drogen.—  No  compounds  of  bromine  and  oxygen  have 
been  prepared  in  a  pure  condition.     With  oxygen  and 
hydrogen  bromine  forms  kypobromous  acid,  bromic  acia, 
and  perbromic  acid. 

The  graphic  formulas  of  these  acids  are  like  those  of  the 
corresponding  chlorine  compounds. 

Hypobromites  are  formed  when  cold,  dilute  alkalies 
are  treated  with  bromine,  but  with  less  than  is  required 
for  saturation. 

2  KOH  +  2  Br »  KOBr  +  H2O  +  KB:-. 


268  FLUORINE,    BROMINE,    IODINE. 

The  hypobromites,  like  the  hypochlorites,  are  oxidizing 
agents. 

Bromates  are  formed  when  hot  alkalies  are  saturated 
with  bromine  (ef.  chlorates,  §  281).  Potassium  bromate 
is  a  white,  crystalline  solid  like  potassium  chlorate. 
Heat  decomposes  it  into  potassium  bromide  and  oxygen. 

2KBrO3=2KBr+3  O2. 

284.  Compounds  of  Iodine  with  Oxygen  and  Hydro- 
gen. —  Two  oxy-acids  of  iodine  are  known  ;  they  are 
iodic  and  periodic  acids. 

lodic  acid  (HIO3  or,  graphically,  H-O-I^^  J  is  formed 

by  oxidizing  iodine  with  concentrated  nitric  acid.  It 
is  a  crystalline  solid.  At  170°  C.  it  breaks  up  into 
iodine  pentoxide  (I2O5)  and  water. 

lodates  are  formed  by  adding  iodine  to  hot,  concen- 
trated solutions  of  alkalies.  With  potassium  hydroxide 
the  equation  is,  —  - 

6  KOH  +61  -  »  KIO3  +  5  KT  +  3 


Periodic  acid  is  HIO,,  or  H-O-I=O.     Its    salts  are 

\> 

periodates,  e.  g.,  sodium  periodate,  NaIO4. 

285.  The  Halogen  Family.  —  From  the  preceding 
pages  it  is  evident  that  there  is  a  great  similarity  in  the 
properties  of  the  elements  fluorine,  chlorine,  bromine, 
and  iodine.  The  close  relation  of  these  elements  to  one 


THE  HALOGEN'  FAMILY.  269 

another  is  yet  more  marked  if  we  consider  that  there 
is  a  gradation  in  the  properties  of  these  elements  in  the 
order  of  the  atomic  masses.  Thus,  the  melting  tempera- 
tures, the  boiling  temperatures,  and  the  specific  gravities 
of  these  elements  rise  from  fluorine  (atomic  mass  19) 
to  iodine  (atomic  mass  127).  The  intensity  of  the 
color  of  these  elements  also  increases  with  the  atomic 
mass :  fluorine  is  greenish  yellow ;  chlorine,  green ; 
bromine,  brown;  and  iodine,  black. 

The  gradation  observed  in  their  physical  properties  is  true 
also  of  their  chemical  properties.  Thus,  the  elements  of  lower 
atomic  mass  can  expel  those  of  higher  atomic  mass  from  their 
soluble  metal  salts. 

The  same  gradation  of  properties  is  noticed  in  the 
compounds  of  the  halogens.  Thus,  the  specific  gravity 
of  the  hydrogen  compounds  increases  from  hydrofluoric 
acid  to  hydriodic  acid ;  while  the  stability  of  these  com- 
pounds decreases  in  the  same  order. 

In  the  case  of  the  compounds  of  the  halogens  with  oxygen, 
and  with  oxygen  and  hydrogen,  the  order  of  stability  is  re- 
versed, iodine  forming  the  most  stable  ones,  chlorine  very  un- 
stable ones,  and  fluorine  none  at  all. 

The  same  gradation  of  color  see^i  in  the  elements 
themselves  appears  in  many  of  their  compounds.  Thus, 
silver  chloride  is  white;  silver  bromide,  light  yellow; 
silver  iodide,  bright  yellow. 

Some  of  the  above  (and  other)  facts  appear  in  the 
following  table :  — 


270 


FLUORINE,    BROMINE,    IODINE. 


PROPERTIES. 

FLUORINE. 

CHLORINE. 

BROMINE. 

IODINE. 

Atomic  Mass  .... 

19 

35.5 

80 

127 

Boiling  Temperature 

—187°  C. 

—33° 

+59' 

•4-184" 

Specific  Gravity   .  . 

1.15  (liquid) 

1.5  (liquid) 

3.2  (liquid) 

5  (solid) 

Union    with    Hydro- 
gen takes  place   . 

In  the  dark 
at  ordinary 
tempera- 
tures. 

In  sunlight. 

At  red  heat. 

A  trod  heat, 
but  incom- 
pletely. 

Pleat  of  formation  of 
Hydrogen  Com- 

37.6 
heat  units. 

22 

8 

—6.1 

Stability  of   Hydro- 
gen Compound  .  . 

Most 
stable. 

Decomposed 
at  1500*  C. 

Decomposed 
at  800°  C. 

Decomposed 
at  180°  C. 

Stability  of  Oxygen 
Compound  .... 

Forms 
none. 

Unstable. 

Forms 
none. 

Most 
stable. 

A  group  of  elements  related  to  one  another  like  the 
halogens  is  called  a  "  Natural  Family  of  Elements/' 
Several  other  natural  families  exist,  and  will  be  referred 
to  later. 


286.  Exercises. 

1.  How  many  grams  of  bromine  can  be  made  from  150  grams 
of  potassium  bromide  by  using  manganese  dioxide  and  dilute 
sulphuric  acid  ? 

2.  Calculate  the  per  cent  of  hydrogen  in  hydriodic  acid.     In 
hydrobromic  acid. 

3.  How  would  you  separate  a  mixture  of  iodine  and  sand  ? 


EXERCISES.  271 

4.  How  could  you  distinguish,  by  chemical  means,  between 
a  chloride,  a  bromide,  and  an  iodide  ? 

5.  How  could  you  identify  a  fluoride,  e.  g.  calcium  fluoride  ? 

6.  About  how  much  would  a  liter  of  air  weigh  at  -f-273°  c. 
and  760  mm.  pressure  ?     A  liter  of  iodine  vapor  ? 


CHAPTER   XVIII. 
OZONE   AND   HYDROGEN  PEROXIDE. 

287.  Ozone. — Oxygen  which -has  been  exposed  to 
the  silent  electric  discharge  possesses  new  properties. 
It  has  a  peculiar  odor  and  oxidizing  powers  not  pos- 
sessed by  ordinary  oxygen.  Thus,  it  oxidizes  silver 
and  mercury  at  once,  whereas  these  metals  are  not  acted 
upon  by  oxygen  at  ordinary  temperatures. 

These  new  properties  are  due  to  the  presence  in  the 
oxygen  of  another  substance,  called  ozone.  The  name 
ozone  is  from  the  Greek  ozein,  to  smell. 

The  same  substance  is  produced  in  almost  every  case  of 
oxidation.  An  illustration  of  this  is  the  slow  oxidation  of 
phosphorus.  If  moist  phosphorus  is  placed  in  a  covered  vessel, 
the  peculiar  odor  of  ozone  soon  appears.  Ozone  is  formed, 
also,  in  the  electrolysis  of  water,  and  appears  at  the  -f-  elec- 
trode along  with  oxygen. 

When  ozone  is  examined,  it  is  found  to  contain 
nothing  but  oxygen ;  it  is,  in  fact,  an  allotropic  form  of 
oxygen  (cf.  §  263).  When  oxygen  changes  into  ozone 
there  is  a  contraction  of  volume  amounting  to  one- 
third  of  the  volume  of  oxygen  taken.  Ozone  is,  there- 
fore, one  and  one-half  times  as  heavy  as  oxygen.  Its 
molecular  mass  is  48  instead  of  32 ;  hence  the  molecule 
of  ozone  contains  three  oxygen  atoms,  and  is  written  O3. 

272 


PROPERTIES  OF  OZONE.  273 

The  peculiar  instability  of  ozone  is  due  to  the  fact  that  the 
change  from  oxygen  to  ozone  is  accompanied  by  absorption  of 
heat.  In  the  presence  of  a  substance  capable  of  taking  up 
oxygen,  the  ozone  molecule  readily  gives  up  an  atom  of  oxy- 
gen, and  thus  reverts  to  molecular  oxygen,  O  =  O. 

288.  Properties  of  Ozone.  — As  might  be  expected, 
the  oxidizing  power  of  ozone  is  very  great.  Moist 
phosphorus  and  sulphur  are  converted  by  it  into  phos- 
phoric and  sulphuric  acids,  respectively ;  and  ammonia 
is  at  once  oxidized  to  nitric  acid.  Organic  coloring 
substances,  e.  g.  indigo  and  litmus,  are  at  once  decolor- 
ized by  ozone.  The  bleaching  of  fabrics  on  exposure 
to  the  air  is  probably  due  to  the  action  of  ozone  pres- 
ent in  the  air. 

When  ozone  is  heated,  its  molecule  is  decomposed,  and 
ordinary  oxygen  results.  The  reversion  of  ozone  to  oxygen  is 
accompanied  by  an  expansion  of  volume  just  equal  to  the  con- 
traction that  takes  place  when  oxygen  changes  into  ozone. 

Ozone  is  readily  absorbed  by  oil  of  turpentine  ;  hence  the 
amount  of  ozone  formed  in  a  given  volume  of  oxygen  may  be 
determined  by  exposing  the  ozonized  oxygen  to  this  substance. 
Only  about  six  per  cent  of  a  given  amount  of  oxygen  can  be 
converted  into  ozone,  because  the  reverse  change  of  ozone 
into  oxygen  soon  produces  a  condition  of  equilibrium. 

The  presence  of  ozone  in  ozonized  air  is  readily  de- 
tected by  means  of  a  mixture  of  potassium  iodide  and 
starch  paste  —  best  upon  a  piece  of  filter  paper.  The 
ozone  sets  iodine  free,  probably  according  to  the  equa- 
tion, 

2  KI  +  H20  +  03 »  2  KOH  +  2  1+  O2. 


274  OZONE  AND   HYDROGEN  PEROXIDE. 

289.  Hydrogen  Peroxide.  —  Closely  related  to  ozone, 
and  long  confused  with  it,  is  hydrogen  peroxide,  H2O2. 
Hydrogen  peroxide  is  a  colorless  liquid  about  one  and 
one-half  times  as  heavy  as  water ;  it  possesses  remark- 
able oxidizing  and  reducing  powers. 

A  dilute  solution  of  hydrogen  peroxide  may  be  made 
by  adding  barium  peroxide,  BaO2,  to  dilute  hydrochloric 
acid.  The  equation  is,  — 

BaO2  +  2  HC1 »  BaCl2  +  H2O2. 

A  somewhat  better  way  is  to  treat  a  dilute  solution  of  tar- 
taric  acid  with  sodium  peroxide,  Na2O2. 

NajOj  +  H2C4H406  =  Na2C4H406  +H2O2. 

Sodium   peroxide   is   made    (along     with    sodium    monoxide, 
Na2O)  by  burning  sodium  in  air  or  oxygen. 

When  phosphorus  is  partly  immersed  in  water,  it  acts 
upon  the  moist  air  to  form  both  hydrogen  peroxide  and 
ozone.  Both  of  these  substances  are  formed,  also,  by 
holding  a  hydrogen  flame  against  a  piece  of  ice. 

Hydrogen  peroxide  is  formed  in  the  electrolysis  of  water,  if 
oxygen  is  passed  into  the  water  at  the  negative  ( — )  electrode. 
The  oxygen  is  reduced  by  the  nascent  hydrogen  evolved  at  the 
electrode. 

Hydrogen  peroxide  is  found  in  the  air,  and  in  all  rain 
water  and  snow. 

290.  Properties  of  Hydrogen  Peroxide.  — Hydrogen 
peroxide  may  be  obtained  almost  pure  by  distilling  a 
dilute  aqueous  solution  of  it  at  low  pressure.     The  ap- 


PROPERTIES   OF  HYDROGEN  PEROXIDE.       275 

paratus  for  distilling  at  reduced  pressure  is  essentially 
as  shown  in  Fig.  57. 


FIG.  57. 

A  distilling  flask  (A)  is  provided  with  a  thermometer  (B) 
and  a  capillary  tube  (C).  The  capillary  tube  allows  a  very 
small  stream  of  air  to  be  drawn  through  the  apparatus.  The 
distilling  flask  is  connected  air-tight  with  the  condenser  (D) 
and  the  receiver  (E).  The  pressure,  in  millimeters  of  mercury, 
is  indicated  by  the  manometer  (F).  The  air  is  exhausted  at  S 
by  a  water  or  mercury  suction-pump. 

At  26  mm.  pressure  hydrogen  peroxide  boils  at  6 9°  C.; 
under  the  same  pressure  water  boils  at  27°  C. ;  hence  the 
two  substances  can  be  separated  readily.  The  aqueous 
solution  of  hydrogen  peroxide  has  a  bitter  taste  and 
produces  white  spots  upon  the  skin.  It  is  a  powerful 
antiseptic. 

Hydrogen  peroxide  decomposes  readily,  especially  in 
the  presence  of  basic  substances.  The  products  are 
water,  and  oxygen  in  the  nascent  condition ;  hence 
hydrogen  peroxide  is  a  powerful  oxidizing  agent.  It 


276  OZONE  AND  HYDROGEN  PEROXIDE. 

decolorizes  indigo,  litmus,  etc.,  as  ozone  does.  It  at 
once  oxidizes  hydrochloric  acid  to  water  and  chlorine. 

Hydrogen  peroxide  acts  also  as  a  reducing  agent  with 
evolution  of  oxygen.  Thus,  it  reduces  mercuric  oxide  to 
mercury  and  sets  oxygen  free. 

HgO  +  H202 »  Hg  +  02  +  H20. 

One  atom  of  each  oxygen  molecule  (O2)  comes  from 
the  hydrogen  peroxide,  and  the  other  from  the  oxide 
reduced. 

Potassium  permanganate  solution  is  at  once  decolorized  by 
hydrogen  peroxide,  and  potassium  chromate  and  bichromate 
solutions  are  changed  to  a  green  color.  All  of  these  are 
reductions. 

Ozone  and  hydrogen  peroxide  reduce  each  other. 
03  +  H202 >  02  +  H20  +  02. 

Hydrogen  peroxide,  like  ozone,  decomposes  with  evolution 
of  heat ;  this  fact  accounts  for  its  instability. 

The  common  test  for  the  presence  of  hydrogen  peroxide 
in  a  solution  is  to  add  to  the  solution  in  a  test  tuhe 
about  two  or  three  cubic  centimeters  of  ether,  and  then 
one  drop  of  potassium  bichromate  solution.  When  the 
test  tube  is  shaken,  the  layer  of  ether  is  colored  a  beau- 
tiful blue,  if  hydrogen  peroxide  is  present. 

291.  Composition  of  Hydrogen  Peroxide.  —  Hydro- 
gen peroxide  is  composed  of  1.01  parts,  by  weight,  of 
hydrogen  to  every  16  parts  of  oxygen.  Its  molecular 
mass  is  34  ;  hence  its  formula  is  H2O2.  The  graphic 
formula  of  hydrogen  peroxide  is  H-O-Q-H. 


COMPOSITION  OF  HYDROGEN  PEROXIDE.       277 

The  effect  of  the  structure  of  the  molecule,  i.  e.,  the  way  in 
which  the  atoms  are  united,  upon  the  properties  of  substances, 
is  admirably  illustrated  by  the  differences  in  the  behavior 
of  the  two  classes  of  dioxides.  Thus,  while  calcium  peroxide 
(CaO2),  sodium  peroxide  (Na2O2),  and  barium  peroxide  (BaO2) 
give  hydrogen  peroxide  when  treated  with  dilute  acids,  lead 
dioxide  (PbO2)  and  manganese  dioxide  (MnO2)  do  not.  The 
structure  of  all  true  peroxides  is  like  that  of  hydrogen  per- 
oxide. The  graphic  formula  of  sodium  peroxide  is,  therefore, 

Ka  —  O  —  O  —  N"a,  and  that  of  barium  peroxide,  O O.     The 

\/ 
Ba 

graphic  formula  of  the  dioxides  is  different,  that  of  manganese 
dioxide   being,  probably,  Mn^,    and    that    of    lead    dioxide 


CHAPTER   XIX. 

THE  NITROGEN  FAMILY. 
PHOSPHORUS,  ARSENIC,  ANTIMONY,  BISMUTH. 

A.     Phosphorus. 

292.  Occurrence  and  Preparation  of  Phosphorus.  — 

Phosphorus  is  found  in  nature  only  in  the  combined 
form,  chiefly  in  phosphates.  The  most  abundant  phos- 
phate is  calcium  phosphate,  Ca3(PO4)2.  Calcium 
phosphate  exists  in  the  soil,  and  is  taken  up  from  it 
by  plants.  Animals  consume  phosphates  in  their  food. 
The  immediate  source  of  most  phosphorus  is  bone-ash, 
which  contains  about  60^  to  70^/  of  its  weight  of 
calcium  phosphate.  The  present  process  of  making 
phosphorus  is  to  heat  calcium  phosphate  with  charcoal 
and  sand  in  the  electric  furnace. 

We    can   understand    the    chemical   reactions   involved   in 
making  phosphorus  by  considering  them  separately. 

(1)  The    calcium   phosphate    probably   breaks    up    in   the 
presence    of  the   silica  (sand)  into  quiclelinie  and  phosphoi-us 
pentoxide. 

Ca3(P04)2 >  3  CaO  +  P2O5. 

(2)  The  silica  and  quicklime  unite  to -give  calcium  silicate. 

3  CaO  +  3  SiO2  =  3  CaSiO8. 

(3)  The   charcoal   reduces    the    phosphorus    pentoxide    to 
phosphorus. 

278 


PROPERTIES.  279 

P205  +  5  C »  2  P+  5  CO. 

Hence  the  complete  equation  is, — 
Ca3(P04)2  +  3  8i02  +  5  C  =  3  CaSiO3  +  5  CO  +  2  P. 

The  phosphorus  escapes  from  the  furnace  as  a  vapor, 
and  is  collected  under  water.  To  purify  it,  it  is  redis- 
tilled and  pressed  in  the  liquid  state  (under  water) 
through  a  bone-ash  filter.  The  phosphorus  thus  ob- 
tained is  a  white  and  transparent  solid.  About  three 
thousand  tons  of  phosphorus  are  made  every  year. 

293.  Properties.  —  Phosphorus,  like  sulphur,  exists 
in  several  allotropic  forms  with  widely  differing  proper- 
ties. Ordinary  or  yellow  phosphorus  has  a  specific 
gravity  of  about  1.8,  melts  at  about  45°  C.,  and  boils 
at  287°  C.  It  is  insoluble  in  water,  but  dissolves 
readily  in  carbon  disulphide,  CS2. 

Phosphorus  derives  its  name,  which  means  "  bearer  of 
light  "  (cf.  Latin,  lucifer),  from  its  property  of  phosphorescing, 
i.  e.,  glowing,  when  exposed  in  the  dark  to  moist  air  or  other 
gases  containing  oxygen.  This  phenomenon  is  caused  by  slow 
combustion  on  the  surface  of  the  phosphorus. 


Ordinary  phosphorus  ignites  in  air  at  40°  C.,  and 
burns  with  a  hot  flame  to  phosphorus  trioxide  and 
pent  oxide. 


303=2P208. 


4  P  +  5  02  =  2  P206. 

The  spontaneous  ignition  of  finely  divided  phosphorus 
has  already  been  described  (cf.  §  29). 


280  THE  NITROGEN  FAMILY. 

Phosphorus  unites  readily  with  chlorine,  bromine, 
and  iodine  even  at  the  ordinary  temperature.  Two 
compounds  of  phosphorus  and  chlorine  are  possible, 
viz.,  the  trichloride,  PC13,  and  the  pentachloride,  PC15. 
Phosphorus  trichloride  is  a  liquid ;  phosphorus  penta- 
chloride,  a  crystalline  solid. 

2  P  -f  3  C12  =  2  PC18. 
2  P  -f  5  C12  =  2  PC16. 

Yellow  pliosphorus  is  very  poisonous. 

294.  Red  Phosphorus.  —  A  great  difference  exists 
between  yellow  phosphorus  and  the  red  modification. 
Red  phosphorus  is  a  reddish  powder  2.2  times  as  heavy 
as  water,  infusible  at  red  heat,  unable  to  phosphoresce, 
insoluble  in  carbon  disulphide,  and  not  poisonous.  It 
ignites  at  about  260°  C.  in  air.  Red  phosphorus  unites 
with  the  halogens  only  when  heated. 

Red  phosphorus  is  prepared  by  heating  the  ordinary  form  in 
closed  iron  tubes  to  300°  C.  A  small  amount  of  the  yellow 
phosphorus  remains  unchanged  ;  this  is  removed  by  means  of 
carbon  disulphide,  in  which  the  red  variety  is  insoluble. 

When  a  given  amount  of  red  phosphorus  is  burned, 
there  is  much  less  heat  liberated  than  with  an  equal 
amount  of  the  yellow  form ;  the  red  has  therefore  much 
less  energy  than  the  yellow.  This  statement  agrees 
with  the  known  fact  that  when  yellow  phosphorus  is 
changed  into  the  red  there  is  an  evolution  of  heat. 


MATCHES.  281 

.  295.  Molecular  Mass  of  Phosphorus.  —  The  weight 
of  a  liter  of  phosphorus  in  the  vapor  condition  is  almost 
four  times  that  of  a  liter  of  oxygen  at  the  same  temper- 
ature and  pressure ;  consequently  the  molecular  mass  of 
phosphorus,  as  a  vapor,  must  be  about  124,  that  is, 
about  four  times  the  atomic  mass.  This  can  be  due 
only  to  the  fact  that  the  molecule  of  phosphorus  in  the 
gaseous  condition  contains  four  atoms.  The  molecular 
formula  of  phosphorus  vapor  is  thus  written  P4,  just  as 
that  of  oxygen  is  O2. 

296.  Matches.  —  Most  of  the  phosphorus  that  is 
made  is  used  to  tip  matches.  The  ordinary  friction 
match,  as  made  at  present,  consists  of  a  splint  of  wood 
tipped,  first,  with  sulphur,  and  then  with  a  mixture  con- 
taining some  oxidizing  agent,  phosphorus,  and  an  ad- 
hesive substance,  like  glue.  The  oxidizing  agent  may 
be  potassium  nitrate  or  chlorate,  or  the  oxide  of  lead 
known  as  red  lead,  which  has  the  formula  Pb3O4. 

The  chemical  operations  involved  in  lighting  a  match  are 
essentially  as  follows  :  — 

(1)  The   heat  generated  by  rubbing  the    tip  of  the  match 
against  a  rough  surface  causes  the  phosphorus  to  combine  with 
the  oxygen  of  the  oxidizing  agent  in  immediate  contact  with 
it. 

(2)  The  combustion  of  the  phosphorus  causes  the  sulphur 
to  be  raised  to  the  kindling  temperature  of  sulphur. 

(3)  The  burning  of  the  sulphur  raises  the  temperature  of 
the  wood  to  the  kindling  point ;  and  the  match  burns. 

Safety  matches  have  not  the  property  of  being  easily 


282 


THE  NITROGEN  FAMILY. 


ignited  when  rubbed ;  they  require  contact  with  a  spe.- 
cially  prepared  surface.  This  surface  is  usually  on  the 
side  of  the  match  box,  and  consists  of  red  phosphorus 
mixed  with  sand.  The  tip  of  the  safety  match  generally 
contains  antimony  trisulphide  (Sb2S3),  an  oxidizing 
agent,  and  glue. 

297.  Hydrogen  Phosphide  (PH3).  —  Hydrogen  phos- 
phide, orphosphine,  is  a  colorless  gas  which,  as  ordinarily 
made,  is  spontaneously  combustible.  The  common  method 
of  preparing  it  is  to  heat  a  mixture  of  yellow  phos- 
phorus and  a  strong  solution  of  sodium  hydroxide.  The 
equation  is,  — 


»  3  NaH2P02  +  PH8. 
sodium  hypo- 
phosphite. 


Hydrogen 


FIG.  58. 


The  apparatus  (Fig.  58)  consists  of  a  generating  flask  con- 
taining the  phosphorus  and  the  sodium  hydroxide  solution.  The 
stopper  of  the  flask  has  two  holes,  one  for  a  tube  from  a  hy- 
drogen generator  and  the  other  for  a  delivery  tube  ending 


SALTS.  283 

under  water.  The  air  of  the  apparatus  is  first  washed  out  by 
means  of  hydrogen  (or  illuminating  gas)  ;  the  gas  is  then  cut 
off  and  the  flask  is  heated.  The  escaping  phosphine  may 
be  collected  in  a  receiver,  as  shown  in  the  figure,  and  this 
exposed  to  the  air,  or  the  bubbles  of  the  gas  may  be  allowed  to 
escape  through  the  water  directly  into  the  air.  The  material 
of  the  white  smoke  formed  when  phosphine  burns  is  phos- 
phorus  pentoxide,  water,  and  phosphoric  acid. 

The  equation  for  the  combustion  of  phosphine  is,  — 
2  PH3  +  4  02 >  P205  +  3  H20. 

Pure  phosphine  is  not  actually  ignited  until  its  tem- 
perature reaches  100°  C.  As  the  gas  is  ordinarily 
prepared,  however,  it  contains  small  amounts  of  the 
vapor  of  another  phosphide  of  hydrogen  (P2H4),  which 
is  spontaneously  combustible,  and  which,  therefore,  ig- 
nites the  phosphine. 

298.  Phosphonium  Salts. — Phosphine  may  be  re- 
garded as  ammonia,  NH3,  with  its  nitrogen  replaced  by 
phosphorus.  Although  similar  to  ammonia  in  composi- 
tion, phosphine  is  much  less  basic.  The  aqueous  solution 
of  phosphine  is  not  alkaline  at  all ;  the  compound 
PH4OH  can  hardly  be  present  in  the  solution. 

Phosphine  can,  however,  be  made  to  unite  with  the 
halogen  acids.  The  compounds  thus  formed  correspond 
with  the  ammonium  salts  of  the  halogens  ;  hence  they 
are  called  phosphonium  salts.  Phospkonium  bromide 
(PH4Br)  and  phosphonium  iodide  (PH4I)  are  much  like 
ammonium  bromide  and  iodide,  respectively. 


284  THE  NITROGEN  FAMILY. 

Phosphonium  iodide  is  decomposed  by  soluble  hydroxides, 
much  as  ammonium  chloride  is  (c/.  §  142).  This  fact  is  evi- 
dent from  the  equations, 


NH4C1  +  KOH  -  »  KC1  -f  NH4OII  (i.  e.,  NH3  -f  II2O)  ; 
PH4I  +  KOH  -  »  KI  +  Pir3  -J-  H2O. 


299.  Phosphides.  —  The  phosphides  of  the  metals 
may  be  considered  derivatives  of  hydrogen  pjiosphide, 
just  as  sulphides  are  of  hydrogen  sulphide.  The  formula 
of  calcium  phosphide  is  Ca3P2  ;  of  silver  phosphide,  AggP. 

Calcium  phosphide  is  a  white  solid  ;  when  it  is  treated  with 
water  or  with  hydrochloric  acid,  it  gives  off  hydrogen  phos- 
phide. The  equation  resembles  that  for  the  action  of  hydro- 
chloric acid  upon  ferrous  sulphide.  Both  equations  are  given. 


Ca3P2  +  G  HC1  -  »  3  CaCl2  +  2  PIT8. 
FeS  +  2  HC1  -  »  Fed,  +  HaS. 

300.  Oxides  of  Phosphorus.  —  Two  common  oxides 
of  phosphorus  are  known,  viz.,  the  trioxide  (P2Og)  and 
the  pentoxide  (P2O5).  Both  are  white  solids. 

The  weight  of  a  given  volume  of  phosphorus  trioxide  in  the 
gaseous  state  is  known  to  be  twice  that  demanded  by  the  formula 
P2O3  ;  consequently  it  is  better  to  write  the  formula  P4O6. 

Phosphorus  pentoxide  is  formed  when  phosphorus 
burns  in  air  or  oxygen  free  from  moisture. 

;  4P  +  502=2P206. 

It  is  a  bulky,  white  solid  which  has  great  attraction  for 


HYPOPHOSPHOROUS  ACID.  285 

moisture  ;  when  put  into  water  it  hisses  like  hot  iron. 
The  product  is  metaphosphorie  acid,  HPO3. 


Phosphorus  pentoxide  has  been  referred  to  as  capable  of 
decomposing  anhydrous  nitric  acid  and  producing  nitrogen 
pentoxide  (cf.  §  166).  It  is  the  anhydride  of  metaphosphorie 
acid,  as  nitrogen  pentoxide  is  of  nitric  acid. 

301.  Oxygen  Acids  of  Phosphorus.  —  Several  com- 
pounds of  phosphorus  with  oxygen  and  hydrogen  are 
known.     Three  of  these  form  a  series  like  the  oxygen 
acids  of  chlorine  (e/.  §§  106  and  278);  they  are,  - 

Hypophosphorous  acid,  HgPO2  ; 
Phosphorous  acid,  H3PO3; 
Phosphoric  acid,  H3PO4. 

302.  Hypophosphorous    Acid.  —  Attention    has    al- 
ready been  called    to    the  fact   that  when    phosphorus 
acts  upon    sodium  hydroxide   (cf.   §   297)   it   produces 
sodium     hypophosphite    (NaH2PO2)    as    well    as     phos- 
phine.     With  barium  hydroxide,  barium  hypophosphite, 
Ba(H2PO2)2,  is  produced. 

The  hypophosphites  are  salts  of  hypophosphorous 
acid.  This  is  a  monobasic  acid  (cf.  §  105).  Its  graphic 

H\   ^0 

formula  is          P^.         ,    only  the    hydrogen     atom    at- 
^ 


tached  to  oxygen  being  ordinarily  replaceable  by  metals. 


286  THE  NITROGEN   FAMILY. 

HypophosphorouB  acid  is  a  powerful  reducing  ayent,  owing 
to  the  ease  with  which  it  goes  over  into  phosphoric  acid. 


303.  Phosphorous  Acid.  —  Phosphorous  acid  is  an 
intermediate  product  in  the  oxidation  of  hypophos- 
phorous  acid.  It  is  itself  a  reducing  agent,  owing  to  its 
ready  oxidation  to  phosphoric  acid.  Its  anhydride  is 
phosphorus  trioxide,  P2O3. 

P203  +  3  H20  =  2  H3P03. 

Phosphorous  acid  may  be  prepared  by  treating  phos- 
phorus trichloride  or  tribromide  with  water  (<;f.  §  271). 

PC18  +  3  H2O  -  »  H8PO8  +  3  11C1. 

Ordinary  phosphorous  acid  is  dibasic  /  its  graphic  formula 

.0 
is,  therefore,  II  —  P  —  OH. 


304.  The  Phosphoric  Acids.  —  The  three  phosphoric 
acids  are,  (1)  orthophosphoric  acid,  or  simply  phosphoric 
acid  (H3PO4);  (2)  pyrophosphoric  acid  (H4P2O7),  and 
(3)  metaphosphoric  acid  (HPO3). 

There  is  still  another  phosphoric  acid,  from  which  all 
of  the  three  named  may  be  supposed  to  be  derived,  />// 
loss  of  water  ;  this  is  normal  phosphoric  acid,  P(OH)r/ 
It  corresponds  to  normal  nitric  acid,  N(OH)5.  But 
while  the  ordinary  form  of  nitric  acid  is  HNO3,  i.  e.,  the 
molecule  of  the  normal  acid  minus  two  molecules  of 
Water,  the  phosphoric  acid  from  which  the  phosphates 


PllEPAliATlOy   OF    THE  PUOSPUOIHC  AC  ID  8.     287 

are  chiefly  derived  is  the  orthophosphoric  acid,  H3PO4. 
The  molecule  of  this  acid  is  the  molecule  of  the  normal 
acid  minus  only  one  molecule  of  water. 

P(OH).-H20  =  H8P04. 

If  the  molecule  of  orthophosphoric  acid  loses  a  mole- 
cule of  water,  we  have  metaphosphoric  acid,  HPO3. 
Tins  acid  corresponds  to  nitric  acid,  HNO3. 

Pyrophosphoric  acid  is  derived  from  orthophosphoric 
acid  by  loss  of  one  molecule  of  water  from  two  molecules 
of  the  orthophosphoric  acid. 

2H,P04  —  H,0  =  H4P,0T. 

The  complete  anhydride  of  all  the  phosphoric  acids  is  phos- 
phorus pentoxide,  P2O6. 

Orthophosphoric  acid  is   tribasic,  a  fact  expressed  in 

^,OH 
its  graphic  formula,  O—  P  —  OH.      It  therefore  forms 


two  acid  salts  (cf.  §  103)  and  a  normal  salt 

Thus,  with  sodium  hydroxide  we  ma}'  get,-  — 

(1)  Sodium  di-hydrogen  phosphate,  N"aH2PO4  ; 

(2)  Disodium  hydrogen  phosphate,  N"a2HPO4  ; 
(tt)  Trisodium  phosphate,  Na3PO4. 

Salts  like  the  first  of  these,  in  which  only  one-third  of  the 
hydrogen  is  replaced,  are  called  primary  phosphates  ;  those  in 
which  two-thirds  of  the  hydrogen  is  replaced  are  called  secondary 
phosphates.  The  normal  salts  are  tertiary  phosphates. 

305.  Preparation  of  the  Phosphoric  Acids.  —  The 

best  way  to  obtain  orthophosphoric  acid  is  to  treat  red 


288  THE  NITROGEN  FAMILY. 

phosphorus  with  nitric  acid,  and  to  evaporate  the  result- 
ing solution.  At  the  ordinary  temperature  the  acid 
consists  of  colorless,  deliquescent  crystals. 

Pyropliosphoric  add  is  best  made  by  heating  the  ortho-  acid 
for  some  time  to  260°  C.  ;  the  meta-  acid  is  made  by  heating 
the  ortho-  or  the  pyro-  acid  to  300°  C. 

Metaphosphorie  acid  always  results  when  phosphorus 
pentoxide  dissolves  in  water. 

P206  +  H20  =  2  HP03. 

When  the  meta-  acid  is  boiled  with  water  it  goes  over 
into  the  ortho-  acid. 


306.  Salts  of  the  Phosphoric  Acids.  —  MetapJios- 
phates  are  obtained  from  primary  orthophosphates  only, 
by  the  loss  of  one  molecule  of  water  from  every  mole- 
cule of  the  orthophosphate. 

Thus,  sodium  metaphosphate,  NaPO3,  is  obtained  by 
heating  sodium  di-hydrogen  phosphate. 

NaH2PO4  =  H2O  -f  NaPO3. 

The  so-called  "  metaphosphate  bead''''  is  made  by  heating 
either  sodium  di-hydrogen  phosphate  or  sodium  ammonium 
hydrogen  phosphate  upon  a  loop  of  platinum  wire.  Sodium 
ammonium  hydrogen  phosphate  (also  called  "  microcosmic 
salt")  has  the  formula  NaNH4HPO4.  When  heated  it  first 
loses  ammonia,  like  any  ammonium  salt  of  a  "fixed"  acid, 
giving  sodium  di-hydrogen  phosphate.  This  then  loses  water. 

+  H2O  . 


USES   OF  THE  PHOSPHATES.  289 

When  a  secondary  orthophosphate,  e.  g.,  Na2HPO4, 
loses  water,  one  molecule  of  water  must  come  from  two 
molecules  of  the  phosphate ;  hence  a  pyrophosphate 
results. 

2  Ka2HPO4  —  H2O  =  Na4P2O7  (sodium  pyrophosphate). 

307.  Uses  of  the  Phosphates.  —  A  knowledge  of 
the  relations  between  the  phosphates  is  essential  to  an 
understanding  of  the  reactions  involved  in  making 
fertilizers  and  phosphorus. 

The  phosphate  found  in  bone-ash  and  in  nature  is 
normal  calcium  phosphate,  Ca3(PO4)2.  This,  however, 
is  insoluble  in  water.  To  convert  it  into  soluble  form 
for  the  use  of  plants,  the  normal  phosphate  is  treated 
with  sulphuric  acid.  This  changes  it  into  primary 
calcium  phosphate,  Ca(H2PO4)2,  which  is  soluble. 

Ca3(P04)2  +  2  H2S04 »  2.  CaSO4  +  Ca(H2PO4)2. 

Calcium  sulphate  is  much  less  soluble  in  water  than  the 
primary  calcium  phosphate  ;  hence  the  two  can  be  separated 
readily. 

The  primary  phosphate  of  calcium  is  used  not  only 
as  a  fertilizer,  but  also  in  making  baking  powders  (cf. 
§  206)  and  as  a  source  of  phosphorus. 

The  old  process  of  making  phosphorus  from  a  phosphate 
consists,  (1)  in  changing  the  phosphate  into  metaphosphate,  and 
(2)  in  reducing  the  metaphosphate. 


290  THE  NIT  HOG  EN  FAMILY. 

The  change  of  primary  calcium  phosphate,  like  that  of  the 
sodium  salt,  into  metaphosphate,  takes  place  on  heating. 

Ca(H2P04)2  =  Ca(P03)2  +  2  H2O. 

The  reduction  of  calcium  metaphosphate  to  phosphorus 
takes  place  when  the  metaphosphate  is  heated  with  charcoal  or 
with  charcoal  and  silica,  SiO2  (cf.  §  292). 


B.  Arsenic. 

308.  Occurrence   and    Preparation   of    Arsenic.  — 

The  element  arsenic  is  found  in  nature  both  free  and 
combined.  Its  chief  ores  are  realgar  and  orpiment 
(As2S2  and  As2S3,  respectively),  the  oxide  (As2O3),  and 
arsenopyrite  (FeAsS).  Arsenopyrite  is  iron  pyrites 
(FeS2)  with  half  of  the  sulphur  replaced  by  arsenic. 

The  element  may  be  prepared  by  reducing  the  oxide 
with  charcoal. 

As2O3  +  3  C »  2  As  +  3  CO. 

309.  Properties   of  Arsenic.  —  Arsenic  forms   com- 
pounds which  correspond  closely  with  the  compounds 
of  phosphorus.     The  element  itself  is,  however,  more 
metallic  than  phosphorus.     It  exists  in  at  least  two  allo- 
tropic  forms. 

The  ordinary  form  of  arsenic  is  gray,  has  a  crystal- 
line structure,  and  is  about  5.7  times  as  heavy  as  water. 
It  is  not  at  all  malleable,  but,  on  the  contrary,  crumbles 
to  powder  when  struck. 


PROPERTIED   OF  ARSENIC.  291 

When  arsenic  is  heated  to  about  500°  C.  out  of  contact  with 
the  air,  it  sublimes,  forming  a  yellow  vapor.  By  comparing  the 
weight  of  a  known  volume  of  this  vapor  with  that  of  oxygen 
under  the  same  conditions,  it  is  found  that  the  molecular  mass 
of  arsenic  vapor  is  about  300.  The  atomic  mass  being  75,  the 
molecule  must  contain  four  atoms  ;  hence  the  molecular  formula 
is  As4.  Above  1,700°  C.,  however,  most  of  the  molecules  con- 
taining four  atoms  dissociate  into  simpler  molecules  of  two 
atoms  each,  i.  e.,  As2  molecules. 

Arsenic  begins  to  burn  at  about  180°  C.  to  form 
arsenic  trioxide,  As2O3.  The  flame  is  bluish-white. 
Like  phosphorus  and  antimony,  arsenic  unites  with 
chlorine  at  the  ordinary  temperature  to  form  the 
chloride,  AsCl3. 

2  As  +  3  C12  =  2  AsCl3. 

The  same  substance  is  formed  when  arsenic  trioxide, 
As2O3,  is  treated  with  concentrated  hydrochloric  acid 
solution. 

As2O3  +  6  HC1 »  2  AsCl3  +  3  H2O. 

Arsenic  trichloride  is  a  colorless  liquid ;  it  is  decom- 
posed by  an  excess  of  water,  giving  arsenious  acid  and 
hydrochloric  acid. 

Cl       HOH  OH 

As  —  Cl  +  HOH »  As  —  OH  +  3  HC1. 

X  Cl       HOH  x  OH 

Arsenic  trichloride  is  thus  like  phosphorus  trichloride, 
which  with  water  gives  phosphorous  acid  and  hydro- 
chloric acid  (ef.  §  303).  ' 


292 


THE  NITBOGEN  FAMILY. 


310.  Hydrogen  Arsenide.  —  Arsenic  combines  with 
nascent  hydrogen  to  form  hydrogen  arsenide  or  ursine, 
AsH3,  a  substance  which  corresponds  with  phosphine 
gas,  PHg.  Arsirie  cannot,  however,  be  made  to  unite 
with  hydrobromic  acid,  etc.,  to  give  compounds  resembling 
ammonium  and  phosphonium  salts  (^f.  §  298). 

The  most  common  method  of  getting  arsine  (mixed 
with  hydrogen)  is  to  add  an  arsenic  compound  to  a  flask 
in  which  hydrogen  is  being  generated;  the  nascent 
hydrogen  unites  with  the  arsenic  of  the  compound. 

Marsh's  Test. — Advantage  is  taken  of  the  properties  of 
arsine  to  test  for  the  presence  of  arsenic  in  any  substance  ;  the 
test  is  known  as  Marsh's  test. 

To  a  flask  in  which  pure  hydrogen  is  being  generated  (Fig. 
59),  there  is  attached  a  calcium  chloride  tube  and  a  hard  glass 


J 


FIG.  59. 


tube  drawn  out  as  shown  in  the  figure.  The  hydrogen  is 
allowed  to  pass  off  until  the  usual  test  ((/.  p.  13)  shows  that 
all  air  has  been  removed.  The  jet  of  hydrogen  is  now  lighted, 


AESENIC   TRIOXIDE.  293 

and,  a  few  drops  of  the  liquid  to  be  tested  for  arsenic  are  added 
through  the  thistle-tube.  If  arsenic  is  present,  the  flame 
changes  to  a  bluish-white  color,  and  a  cold  piece  of  porcelain 
held  in  the  flame  receives  a  shiny,  black  deposit,  called  an 
"  arsenic  mirror." 

If  the  hard  glass  tube  is  heated,  the  arsine  passing  through 
it  is  decomposed,  and  an  arsenic  mirror  appears  in  the  tube. 
Here  the  arsenic  may  be  identified  by  passing  hydrogen  sul- 
phide, H2S,  through  the  heated  tube.  The  same  precautions 
must  be  taken  to  have  all  air  removed  as  in  the  case  of  hydro- 
gen. Hydrogen  sulphide  changes  the  arsenic  into  arsenic 
trisulphide,  As2S3 ;  this  is  a  golden-yellow  solid  called  orpiment 
(from  auri  pigmentum).  If,  now,  dry  hydrochloric  acid  gas  is 
passed  through  the  tube,  the  arsenic  trisulphide  does  not 
change.  These  properties  serve  to  distinguish  between  the 
arsenic  mirror  and  that  of  antimony  (cf.  §  318). 

311.  Arsenic  Trioxide.  —  The  oxides  of  arsenic  are 
the  trioxide  (As2O3)  and  the  pentoxide  (As2O5) ;  these 
correspond  to  the  phosphorus  oxides.  Arsenic  trioxide 
(often  called  "arsenic"  or  "white  arsenic")  is  the  most 
common  arsenic  compound.  It  is  a  white  powder  which 
sublimes,  without  melting,  at  about  220°  C.  The  vapor 
has  a  garlic  odor.  When  the  vapor  solidifies,  the  arsenic 
trioxide  appears  in  the  form  of  a  transparent  mass. 
At  very  high  temperatures  the  molecule  of  the  vapor 
is  represented  by  As2O3;  but  between  220°  and  700°  C. 
the  molecules  are  doubled,  and  the  formula  is  As4O6. 

Uses.  —  Arsenic  trioxide  is  used  in  medicine  and  as  a 
poison.  Its  poisonous  action  upon  the  human  system 
is  rather  slow,  owing  to  its  dissolving  only  slowly  in  the 


294  THE  NITROGEN  FAMILY. 

liquid  of  the  stomach.  The  antidote  is  a  mixture  of 
ferric  hydroxide  [Fe(OH)3]  and  magnesia  (MgO). 

The  people  of  certain  mountain  districts  are  addicted  to  the 
use  of  arsenic  trioxide  because  it  enables  them  to  breathe  more 
easily  when  climbing.  By  beginning  with  very  small  quanti- 
ties and  gradually  increasing  the  dose,  they  are  able  to  take 
much  more  than  the  lethal  dose  without  injury.  But  the  diffi- 
culty comes  when  they  try  to  leave  off  the  habit  ;  for  they  then 
suffer  all  the  effects  of  arsenic  poisoning. 

312.  Arsenious  Acid.  —  Arsenic  trioxide  is  slightly 
soluble  in  water;  the  solution  probably  contains  arseni- 
ous  acid,  HoAsOq. 

o  rf 

As.2O3  +  3  H2O  =  2  H3AsO3. 

Arsenious  acid  is  not  known  in  the  free  state  because  it 
breaks  up  into  arsenic  trioxide  (its  anhydride)  and 
water.  The  salts  of  arsenious  acid  are  called  arsenites. 
Solutions  of  these  are  formed  when  arsenic  trioxide  is 
treated  with  alkalies.  Thus,  sodium  hydroxide  and 
arsenic  trioxide  (with  water  this  is  arsenious  acid)  form 
sodium  arsenite,  Na3As()3. 

H3AsO3  +  3  NaOH  -  »  Na3AsO3  +  3  H2O. 

Many  arsenites  are  derived  from  nwtarsenious  acid,  HAsO2, 
which  may  be  looked  upon  as  arsenious  acid  minus  water. 

IT3AsO3  =  TT  As(X  +  II2O. 
Sodium  metar  senile  would  be 


Double    Nature   of   Arsenious  Acid.  —  Arsenic   tri- 
oxide (or  arsenious  acid)  reacts  not  only  with  alkalies, 


ARSENIC  ACID.  295 

giving  arsenites  and  water,  but  also  with  acids,  giving 
arsenic  salts  and  water.  Thus,  with  concentrated  hy- 
drochloric acid  and  arsenic  trioxide,  we  get  arsenic  tri- 
chloride and  water. 

As2O3  +  6  HC1 »  2  AsCl3  +  3  H2O ;  or, 

H3AsO3  +  3  HC1 >  AsCl3  +  3  H2O. 

Arsenious  acid  has  thus  a  double  nature ;  for  toward 
strong  bases  it  acts  like  an  acid,  forming  with  the  base 
an  arsenite;  while  toward  a  strong  acid  it  acts  like  a 
base,  giving  with  the  acid  a  salt  and  water.  Arsenic  is, 
in  fact,  intermediate  between  the  non-metals  and  the 
metals. 

Arsenic  Greens.  —  At  least  two  arsenic  compounds  have 
a  bright  green  color  ;  these  are  c-oppe*  arsenite,  called  "  Scheele's 
green"  and  a  mixture  of  copper  arsenite  and  copper  acetate, 
called  "  Schweinfurt's  green."  Both  of  these  are  sold  as 
ki  Paris  green."  These  dyes  were  formerly  used  to  color  wall- 
paper, paints,  etc.  They  are  too  dangerous,  however,  and  are 
now  rarely  used  as  dyes.  Paris  green  is  used  to  destroy  potato- 
bugs,  etc. 

313.  Arsenic  Acid.  —  Arsenic  pentoxide,  As2O6,  is 
the  anhydride  of  arsenic  acid,  H3AsO4. 

As205  +  3  H20  =  2  H3As04. 

Arsenic  acid  is  formed  by  dissolving  arsenic  or  arsenic 
trioxide  in  concentrated  nitric  acid.  (Compare  the 
preparation  of  phosphoric  acid  from  phosphorus,  §  305.) 


296  THE  NITROGEN  FAMILY. 

The  arsenic  acids  have  formulas  of  the  same  type  as 
those  of  phosphorus  :  — 

H3AsO4  is  orthoarsenic  acid  ; 
H4As2O7  is  pyroarsenic  acid  ; 
HAsO3  is  metarsenic  acid. 

Metarsenic  acid  finally  gives,  by  loss  of  water,  arsenic  pentoxide. 
2  HAsO3  =  H2O  +  As2O5. 

The  arsenates  are  like  the  corresponding  phosphates, 


314.  Arsenic  Trisulphide.  —  When  the  solution  of 
an  arsenic  compound  is  treated  with  hydrogen  sulphide, 
a  yellow  precipitate  is  generally  produced  ;  this  consists 
of  either  the  trisulphide  (As2S3)  or  the  pentasulphido 
(As2S5).  Both  sulphides  react  with  ammonium  sul- 
phide [(NH4)2S]  and  other  soluble  sulphides.  The  so- 
lution contains  a  sulpharsenite.  or  sulphar  senate.  Thus 
with  sodium  sulphide  and  arsenic  trisulphide  the  equa- 
tion is,  — 

3  Na2S  +  As2S3  =  2  Na3AsS3. 

The  sulpharsenite  (Na3AsS3)  is  simply  an  arsenite  with 
its  oxygen  replaced  by  sulphur. 

When  the  sulpharsenite  is  treated  with  a  dilute  acid,  e.  g.. 
hydrochloric  acid,  sulphar  senious  acid,  H3AsS3,  is  set  free  ;  this 
breaks  up  into  hydrogen  sulphide  and  arsenic  trisulphide. 
Arsenic  trisulphide,  being  insoluble,  is  reprecipitated.  These 
facts  are  shown  in  the  equations,  — 


PHYSICAL   PROPERTIES.  297 


+  3  HC1  -  »  H3  AsS,  +  3  NaCl. 
2  H3AsS3  =  3  H2S  -f  As2S3. 

Ammonium   sulpharsenite,   (NTf4)3AsS3,  undergoes    a  similar 
decomposition. 

C.  Antimony. 

315.  Preparation  of  Antimony.  —  Antimony  is  found 
in  nature  chiefly  combined  with  sulphur  in  the  mineral 
stibnite,  Sb2S3.  To  obtain  the  element  the  sulphide  is 
first  roasted,  i.  e.,.  heated  in  a  stream  of  air,  and  then 
heated  with  charcoal. 

Roasting  converts  the  antimony  sulphide  into  the  tnoxide 
(Sb2O3),  or  the  tetroxide  (Sb2O4),  and  sulphur  dioxide. 

2  Sb2S3  +  9  02  -  »  2  Sb203  +  G  SO2  ;  or, 
Sb2S3  -f  5  O2  -  -  »  Sb2O4  -f  3  SO2. 

The  reduction  of  either  of  these  oxides  by  charcoal  gives  anti- 
mony and  carbon  monoxide. 

Sb2O3  +  30  -  »  2  Sb  +  3  CO. 


316.  Physical  Properties.  —  Antimony  is  a  solid 
having  a  bright,  silvery  luster  which  is  not  easily  tar- 
nished in  air.  Antimony  is  not  malleable.  At  about 
430°  C.  it  melts  to  a  liquid  of  a  slightly  higher  specific 
gravity.  When  melted  antimony  solidifies  it  expands 
again;  hence  antimony  is  valuable  as  a  constituent  of 
materials  for  casts,  such  as  type-metal.  The  specific 
gravity  of.  the  solid  is  6.7.  The  specific  gravity  of  the 
yapor  shows  that  in  the  gaseous  condition  the  formula  of 
the  molecule  is  Sb2  (<?/,  §§  295  and  300), 


298  THE  NITROGEN  FAMILY. 

317.  Chemical  Properties.  — Antimony  burns  in  the 
air  to  form  the  trioxide  or  the  tetroxide.  It  combines 
with  chlorine  to  give  antimony  trichloride  (ef.  §  84)  or 
the  pvntachloride,  SbCl5.  With  fluorine,  bromine,  and 
iodine  it  forms  similar  compounds. 

Antimony  is  insoluble  in  hydrochloric  acid.  Nitric  acid 
oxidizes  it  to  antimony  trioxide  or  antimonic  acid,  H3SbO4. 

4  Sb  +  3  O2  =  2  Sb2O3. 
4  Sb  +  5  02  =  2  Sba06. 
Sb2O5  +  3  H2O  =  2  II3SbO4. 

With  aqua  regia  antimony  reacts,  giving  antimony 
trichloride.  When  the  solution  is  distilled  it  gives  the 
trichloride  as  a  liquid  boiling  at  223°  C.  This  solidifies 
to  a  pasty  mass  called  "butter  of  antimony." 

Concentrated  sulphuric  acid  reacts  with  antimony. 
The  products  are  shown  in  the  equation, 

2  Sb  +  6  H2S04 »  Sb2(S04)3  +  3  H2SO3  +  3  II2O. 

antimony  sulphate 

Antimony  resembles  arsenic  and  phosphorus  on  the 
one  hand  and  bismuth  on  the  other.  It  is  like  the 
former  elements  in  the  general  structure  of  its  com- 
pounds ;  like  the  latter  in  its  ability  to  form  a  salt  with 
sulphuric  acid  and  in  other  metallic  properties. 

Its  positive  (-}-)  properties  are,  however,  weak,  as  is 
indicated  by  the  easy  decomposition  of  its  salts  by  water. 

Thus,  antimony  trichloride  is  decomposed  by  much  water, 
giving  a  basic  chloride  (its  simplest  formula  is  SbOCl)  and  hy- 
drochloric acid. 


OTHER  ANTIMONY  COMPOUNDS.  299 

Cl       HOH  OH 

Sb  Cl  -f  HOH »  Sb  OH  +  2  HC1. 

Cl  Cl 

OH 
Sb  OH  =  SbOCl  +  H20. 

Cl 

Again,  although  antimony  trioxide  reacts  with  concentrated 
sulphuric  acid,  giving  antimony  sulphate,  Sb2(SO4)3,  yet  when 
antimony  trioxide  reacts  with  the  dilute  acid,  the  product  has 
the  formula  (SbO)2SO4.  The  compound  SbOCl  may  be  called 
antimony  oxy-chloride  or  antimonyl  chloride,  the  group  SbO 
being  called  antimonyl.  The  compound  having  the  formula 
(SbO)2SO4  is,  therefore,  antimonyl  sulphate. 

Antimony,  even  more  than  arsenic,  is  a  transition  ele- 
ment. Its  oxide,  Sb2O3,  reacts  not  only  with  acids, 
giving  antimony  salts,  but  also  with  alkalies,  giving  salts 
of  antimony  acids.  Tims,  antimony  trioxide  gives  with 
sodium  hydroxide  sodium,  antimonite,  Na3SbO3.  This  is 
a  salt  of  antimonious  acid,  which  corresponds  with  phos- 
phorous and  arsenious  acids. 

318.  Other  Antimony  Compounds. — Among  the  im- 
portant antimony  compounds  are  hydrogen  antimonide 
(or  stibine),  tartar  emetic,  and  antimony  trisulphide. 

Hydrogen  antimonide,  SbHg,  is  formed  from  an  anti- 
mony compound  just  as  hydrogen  arsenide  is  formed  from 
an  arsenic  compound,  namely,  by  reduction  with  nascent 
hydrogen  (<?/.  §  310). 

Marsh's  test  may  be  carried  out  with  antimony  in  the  appara- 
tus used  for  arsenic  (Fig.  59).  When,  however,  hydrogen  sul- 
phide is  passed  over  the  antimony  mirror,  the  antimony  sulphide 


300  THE  NITEOGEN  FAMILY. 

formed  is  black.,  while  that  of  arsenic  is  yellow.  Again,  when 
hydrochloric  acid  gas  is  passed  through  the  tube,  the  antimony 
sulphide  forms  drops  of  antimony  trichloride,  while  the  arsenic 
trisulphide  is  unchanged.  We  can  thus  distinguish  between 
arsenic  and  antimony. 

Tartar  emetic  is  potassium  antimony!  tartrate, 
K.SbO.C4H4O6.  It  is  formed  by  heating  a  mixture 
of  antimony  trioxide,  potassium  hydrogen  tartrate 
(cream  of  tartar),  and  water.  It  is  used  in  medicine. 

Antimony  trisulphide,  Sb2S3,  is  formed  by  treating 
solutions  of  either  antimonious  salts  or  antimonites  with 
hydrogen  sulphide.  Two  isomeric  antimony  trisulphides 
are  known.  The  one  formed  by  precipitation  is  brick- 
red,  while  stibnite  is  black.  The  red  form  is  unstable ; 
it  goes  over  into  the  black  form  with  evolution  of  heat. 

The  precipitate  of  antimony  trisulphide  reacts  readily  with 
alkaline  sulphides  giving  sulphantimonites,  just  as  arsenic  tri- 
sulphide gives  sulpharsenites  (cf.  §  314).  The  sulphantimonites 
are  decomposed  by  dilute  acids  ;  and  antimony  trisulphide  is 
reprecipitated. 

319.  Uses   of    Antimony.  —  Antimony    is    used    in 
making   alloys.     Examples    are :     printer's    type-metal, 
pewter,  and  Britannia  metal.     Antimony  black  is  a  finely 
divided  form   of  the   metal;  plaster  casts   are   rubbed 
with  it  to  give  them  a  metallic  coating. 

D.  Bismuth. 

320.  Preparation  of  Bismuth.  —  Bismuth  is  found  in 
nature  free  and  also  in  the  form  of  the  sulphide  (Bi2S3) 


BISMUTH   SALTS.  301 

and  the  oxide  (Bi2O3).  It  is  prepared  from  its  sulphide 
as  antimony  is,  namely,  by  first  roasting  the  sulphide  to 
form  the  oxide,  and  then  reducing  the  oxide. 

To  get  bismuth  free  from  impurities,  such  as  arsenic,  etc., 
it  is  mixed  with  saltpeter  and  heated.  The  impurities  are  thus 
oxidized  to  compounds  soluble  in  water,  and  can  be  separated 
from  the  bismuth. 

321.  Properties.  —  In   its    physical    appearance    bis- 
muth is  much  like  antimony,  but  it  has  a  slightly  reddish 
color.     Its  melting  temperature  is  265°  C.,  and  its  spe- 
cific gravity  about  10. 

Bismuth  burns  in  the  air  when  at  red  heat ;  the  prod- 
uct is  the  trioxide,  Bi2O3. 

Bismuth  trioxide  is  formed,  also,  by  heating  the  nitrate 
Bi(NO3)3  (see  §  169),  and  by  heating  the  hydroxides  Bi(OH)3 
and  BiO.OH,  which  are  produced  when  potassium  hydroxide 
solution  is  added  to  a  solution  of  a  bismuth  salt. 

Unlike  the  trioxides  of  arsenic  and  antimony,  bismuth 
trioxide  has  not  the  ability  to  react  with  alkalies  to  form 
salts.  It  is  an  entirely  basic  oxide.  The  higher  oxide 
of  bismuth,  Bi2O5,  has  slightly  acidic  properties;  for  it 
is  the  anhydride  of  bismuthic  acid,  HBiO3.  The  acid  is, 
however,  very  weak  and  very  unstable. 

Bismuth  forms  no  hydrogen  compound  corresponding 
to  ammonia,  phosphine,  arsine,  and  stibine. 

322.  Bismuth  Salts.  —  Bismuth  combines  with  the 
halogens,  giving    tri-halogen  compounds,  which  corres- 


302  THE  NITEOGEN  FAMILY. 

pond  to  those  of  arsenic  and  antimony.  It  reacts  with 
nitric  acid  to  give  the  nitrate,  Bi(NO3)3;  with  aqua 
regia  to  give  the  chloride,  BiCl3,  and  with  concentrated 
sulphuric  acid  to  give  the  sulphate,  Bi2(SO4)3. 

Bismuth  salts,  like  those  of  antimony,  are  decomposed  by 
much  water,  giving  the  basic  salt  and  free  acid.  The  trichlor- 
ide usually  decomposes  as  follows  :  — 

Cl       HOH  OH 

Bi  Cl  +  HOH »  Bi  OH  -f  2  HC1. 

Cl  Cl 

Bi(OH)2Cl »  BiO.Cl  +  H2O. 

The  nitrate  decomposes  in  a  similar  way. 

When  hydrogen  sulphide  is  added  to  the  solution  of 
a  bismuth  salt  there  is  produced  an  almost  black  pre- 
cipitate of  bismuth  sulphide,  Bi2S3.  This  is  insoluble  in 
alkaline  sulphides. 

323.  Uses  of  Bismuth.  —  The  principal  use  of  bis- 
muth is  as  an  ingredient  of  alloys.  Its  chief  'alloys  are 
Wood's  metal  and  Rose's  metal. 

Wood's  metal  consists  of  four  parts,  by  weight,  of 
bismuth,  one  each  of  tin  and  cadmium,  and  two  of  lead. 
It  melts  at  about  65°  C. ;  hence  it  is  much  used  for 
metal  baths  in  the  laboratory. 

Rose's  metal  contains  nine  parts  of  bismuth  to  one 
each  of  lead  and  tin ;  it  melts  at  94°  C. 

The  basic  nitrate  of  bismuth  (simplest  formula,  BiO.KO3)  is 
used  in  medicine  under  the  name  bismuth  sub-nitrate. 


EXEBCISES.  303 

324.  The    Nitrogen   Family.  —  The  elements  nitro- 
gen, phosphorus,    arsenic,  antimony,   and    bismuth,  to- 
gether with  some  rare  elements,  form  a  natural  family, 
just  as  the  halogens  do ;  for  in  this,  the  nitrogen  family, 
just  as  in  the  case  of  the  halogens,  we  have  a  series  of 
elements  exhibiting  a    gradation    of   properties    in    the 
order  of  the  atomic   masses.      The    table    on  the  next 
page  shows  this  for  some  properties. 

A  complete  list  of  the  properties  of  the  members  of 
the  nitrogen  family  will  not  agree  so  well  as  in  the  case 
of  the  halogens.  When,  however,  we  can  compare 
corresponding  compounds  having  the  same  number  of 
atoms  to  the  molecule,  a  fair  degree  of  regularity  exists. 

325.  Exercises. 

1.  What  means  are  there  of  kindling  fires  without  the  use  of 
phosphorus  ? 

2.  How  much  phosphorus  is  there  in  440  grams  (about  one 
pound)  of  bone-ash,  if  70%  of  the  ash  is  calcium  phosphate, 
Ca3(P04)2? 

3.  Write   the   simplest   equation  for  the    decomposition   of 
phosphonium  bromide,  PH4Br,  by  barium  hydroxide,  Ba(OH)2 
(c/.  §  298). 

4.  Write  the  formulas  of  the  following  substances  :  normal 
'barium  phosphate,  primary   ammonium  phosphate,  potassium 
hypophosphite,  strontium   metaphosphate,  and  silver  pyrophos- 
phate. 

5.  Why  is  it  undesirable  that  arsenic  compounds  should  be 
used  to  color  wall  papers  ?     Explain. 

6.  Show  by  a  simple  equation  the  reduction  of  arsenic  tri- 
oxide  by  nascent  hydrogen  to  arsine. 


304 


THE  NITROGEN  FAMILY. 


7.  When  a  bath  of  Wood's  metal  is  being  melted  the  un- 
melted  metal  floats  upon  the  liquid  portion.  Compare  the 
specific  gravity  of  the  solid  with  that  of  the  liquid.  Will  the 
liquid  expand  or  contract  on  solidifying  ? 


ELEMENT. 

Nitrogen. 

Phosphorus. 

Arsenic. 

Antimony. 

Bismuth. 

ATOMIC 
MASS. 

14 

31 

75 

120 

207 

SPECIFIC 
GRAVITY. 

0.885 
(liquid). 

1.83  to  2.19 

4.7  to  5.7 

6.7 

9.9 

BOILING 

TEMPERA- 

—194° C. 

287° 

450° 

1500° 

1300° 

TURE. 

BOILING 

TEMPERA- 
TURE  OF  TRI- 

Not known, 
but  low. 

76° 

130" 

223C 

447° 

CHLORIDE. 

PROPERTIES 

OF 

TRIOXIDES. 

N80S  is 
anhydride 
of  nitrous 
acid. 

P2O3  is  anhy- 
dride of  phos- 
phorous acid. 
Weaker  than 
nitrous  acid. 

As2O3  is 
both  weak- 
ly acid  and 
weakly 
basic. 

SbjO,,  is 
likeAs2Oa. 

Bi203  has 
only  basic 
properties. 

PROPERTIES 

OF 

PENTOX- 

N,05  is 
anhydride 
of  nitric 
acid. 

P2O5  is  anhy- 
dride of  phos- 
phoric  acid. 
Weaker   than 
nitric  acid. 

As,O5  is 
anhydride 
of  arsenic 
acid. 
Weaker 

Sb,05  is 
anhydride 
of 
antimonic 
acid. 

Bi205  has 
only/oin<- 
ly     acid 
properties. 

IDES. 

than  phos- 
phoric 
acid. 

HYDROGEN 
COM- 
POUNDS. 

NH,,  unites 
with  HC1, 
HI,  etc  ,  to 
form  salts. 

PHS  unites  with 
acids  with  dif- 
ficulty. 

AsH3  does 
not    unite 
with  acids. 

SbH3  does 
not    unite 
with  acids. 

BiH,  is 
not  known 
to  exist. 

CHAPTER   XX. 
THE   PERIODIC   SYSTEM. 

326.  Natural  Families.  —  During  the  first  half  of 
the  nineteenth  century  various  attempts  were  made 
to  classify  the  elements.  Chemists  recognized  the  fact 
that  there  were  "  natural  families "  of  elements,  and 
that  the  members  of  the, same  family  (called  homolog- 
ous elements),  while  bearing  a  general  resemblance  to 
one  another,  yet  showed  a  regular  gradation  of  properties 
in  the  order  of  the  atomic  masses. 

We  have  already  described  two  of  these  families,  viz., 
the  halogen  family  (<?/".§  285)  and  the  nitrogen  family 
(?/'•  §  324). 

Other  natural  families  exist.  Thus,  sulphur  and  oxygen, 
with  the  rarer  elements  selenium  (Se)  and  tellurium  (Te)  form 
a  group  of  homologous  elements.  The  atomic  masses  are,— 

O,  16  ;  S,  32  ;  Se,  79  ;  Te,  127. 

Again,  the  elements  lithium,  sodium,  and  potassium,  with 
the  rare  elements  rubidium  and  ccesmm,  form  the  well-known 
"alkali"  group  of  metals.  The  atomic  masses  of  the  three 
more  common  metals  of  the  family  are, — 

Li,  7  ;  Na,  23  ;  K,  39. 

Here,  as  in  the  case  of  the  halogens,  we  find  a  continuous 
gradation  of  properties  in  the  order  of  the  atomic  masses. 

305 


306  THE  PERIODIC  SYSTEM. 

Lithium  hydroxide,  LiOH,  is  a  weaker  base  than  sodium  hy- 
droxide, NaOH ;  while  potassium  hydroxide,  KOH,  is  the 
strongest  base  of  the  three. 

327.  The  Periodic  Arrangement.  —  Although  the 
connection  between  the  properties  and  the  atomic  masses 
of  elements  in  the  same  group  had  been  recognized  for 
years,  it  was  left  to  the  Russian  chemist  Mendelejeff  and 
the  German  chemist  Lothar  Meyer  to  discover,  in  1869 
to  1871,  a  new  and  peculiar  relation  between  the  proper- 
ties of  all  elements  and  their  atomic  masses.  This  rela- 
tion is  the  basis  of  the  Periodic  System. 

Let  us  write  the  symbols  of  the  first  sixteen  elements  in 
the  order  of  the  atomic  masses.  Omitting  hydrogen,  which 
for  the  present  stands  almost  unrelated,  we  have,  — 

Element,  Li  Gl(Be)  B      C  N"  O   Fl  Na  Mg  Ai 

Atomic  Mass,  7         9        11     12  14  16  19    23    24    27 

Element,  Si        P        S      Cl  K  Ca. 

Atomic  Mass,  28       31       32  35.5  39  40. 

A  study  of  the  elements  from  lithium  to  fluorine,  in- 
clusive, shows  that  there  is  a  regular  gradation  of  prop- 
erties. The  strongly  metallic  properties  of  lithium  are 
weaker  in  glucmum,  and  yet  weaker  in  boron.  The  hy- 
droxide of  boron  is,  in  fact,  called  boric  acid.  In  carbon 
we  have  an  element  with  faintly  electro-negative,  i.  e. 
non-metallic,  properties ;  the  elements  nitrogen  and  oxy- 
gen are  still  more  electro-negative  ;  until  in  fluorine  we 
have  a  typical  non-metal,  and  probably  the  most  electro- 
negative substance  known. 


THE  PERIODIC  ARRANGEMENT.  307 

The  increase  of  atomic  mass  from  7  to  19  has  thus  continu- 
ously diminished  the  electro-positive,  or  metallic,  character 
possessed  by  lithium,  and  has  increased  the  electro-negative 
character  typified  by  fluorine. 

But  after  fluorine  the  gradation  of  properties  does 
not  continue ;  for  the  element  sodium,  the  next  greater 
in  atomic  mass,  is,  like  lithium,  one  of  the  most  typical 
metals.  There  is,  in  fact,  a  sudden  "  reversion  to  type  " ; 
for  sodium  belongs  in  the  same  natural  family  with 
lithium. 

Let  us,  then,  proceed  in  the  order  of  atomic  mass, 
writing  the  second  set  of  seven  elements  under  the  first 
set,  as  in  the  table  in  §  328.  Sodium  will  be  under 
lithium,  magnesium  under  glucinum,  etc.  We  find  the 
same  gradation  of  properties  with  increase  of  atomic 
mass  from  sodium  to  chlorine  as  we  found  from  lithium 
to  jluorine. 

Magnesium,  like  sodium,  is  a  metal ;  but  magnesium  hydrox- 
ide is  a  weaker  base  than  sodium  hydroxide.  A-lurnvnuvfi,  the 
next  in  the  order  of  atomic  mass,  is  also  metallic  ;  but  its  hy- 
droxide is  either  a  base  or  an  acid,  according  to  circumstances. 
In  silicon,  the  next  element,  metallic  properties  are  wanting ; 
silicon  forms  silicic  acid.  Next  come  phosphorus  and  sulphur, 
undoubted  non-metals  ;  and  then  chlorine,  the  first  homologue 
of  fluorine. 

The  element  following  chlorine  is  potassium  ;  its  atomic 
mass  is  39.  Potassium  is  a  typical  metal,  and  belongs 
in  the  same  family  with  lithium  and  sodium.  In  passing 


308  THE  PEBIODIC  SYSTEM. 

from  chlorine  to  potassium  we  have,  therefore,  a  second 
instance  of  "  reversion  to  type." 

The  properties  of  calcium  are  less  metallic  in  character  than 
those  of  potassium  ;  for  calcium  hydroxide  is  a  weaker  base 
than  potassium  hydroxide.  Thus,  after  potassium,  as  after 
sodium,  variation  in  the  properties  of  the  elements  goes  on 
continuously  with  increase  of  atomic  mass  for  another  period. 

From  the  study  of  "  natural  families  "  we  learned  that 
the  properties  of  the  elements  in  any  one  family  vary 
continuously  with  the  atomic  mass ;  now  we  see  that  the 
properties  of  all  elements,  while  not  varying  continu- 
ously, as  in  the  natural  family,  yet  vary  periodically  with 
the  atomic  mass.  That  is  to  say,  a  certain  increase  of 
atomic  mass  is  accompanied  by  a  recurrence  of  certain 
properties  possessed  by  an  element  of  lower  atomic  mass. 
The  facts  are  summed  up  in  what  is  often  called  the 
Periodic  Law,  which  is,  "The  properties  of  the  elements 
are  periodic  functions  of  their  atomic  masses" 

328.  Regularities  in  the  Periodic  Arrangement.  - 
If  we  write   down  the  symbols  of   the  elements   from 
lithium  to  calcium,  putting  similar  elements  in  the  same 
vertical  columns,  we  have  an  arrangement  like  the  fol- 
lowing :  — 

Li  (7)  Gl  or  Be  (9)    B    (11)  C  (12)   N  (14)    O  (16)    Fl  (19) 
Na  (23)      Mg  (24)  Al  (27)    Si  (28)    P  (31)     S  (32)  Cl  (35.5) 
K    (39)       Ca  (40) 

In  this  table  we  observe  several  regularities :  — 


OF  AN  ELEMENT  DETERMINED.  309 


(1)  After  every  period  of  seven  elements  a  second  period  of 
seven  begins. 

(2)  The  difference  between  the  first  and  the  eighth,  the 
second  and  the  ninth,  etc.,  elements  is  in  every  case  nearly  six- 
teen units.     Two  successive  elements  of  the  same  family  are 
thus  separated  by  six  intervening  elements,  and  differ  from  each 
other  by  about  sixteen  units  of  atomic  mass. 

(3)  Elements  in  adjacent  positions  in  the  horizontal  rows, 
i.  e.,  Jieterologous  elements,  differ  from  one  another  by  small 
numbers  of  units,  generally  only  one  or  two.     The  largest  dif- 
ferences occur  between  fluorine  and  sodium  and  between  chlor- 
ine and  potassium,  i.  e.,  at  the  break  in  the  periods. 

329.  Properties  of  an  Element  Determined.  —  We 

may  call  the  two  elements  adjacent  to  another  element 
in  the  horizontal  rows  the  adjacent  Jieterologues  of  the 
element.  Thus,  boron  and  nitrogen  are  the  adjacent  het- 
erologues  of  carbon.  The  two  adjacent  heterologues 
and  the  two  adjacent  Jiomologues  of  an  element  may  be 
called  the  analogues  of  the  element.  Thus,  glucinum, 
sodium,  calcium,  'and  aluminum  are  the  analogues  of 
magnesium. 

Mendelejeff  showed  that  if  the  properties  of  magnes- 
ium were  wholly  unknown  they  could  be  deduced  ap- 
proximately from  the  properties  of  its  four  analogues. 

Thus,  the  atomic  mass  of  magnesium  (24)  is  very  nearly  the 
average  of  the  atomic  masses  of  its  analogues. 

9  +  40  +  23  +  27  =  24  75 

4 

Again,  the  average  of  the  specific  gravities  of  the  analogues  of 
magnesium  gives  approximately  the  specific  gravity  of  magnes- 


310  THE  PERIODIC  SYSTEM. 

ium  itself.  The  specific  gravities  of  sodium,  aluminum,  glu- 
cinum,  and  calcium  are  0.97,  2.56,  2.10,  and  1.58,  respectively. 
The  average  is  1.8.  Experiment  shows  the  specific  gravity  of 
magnesium  to  be  1.75. 

Up  to  the  time  of  Mendelejeif  and  Meyer,  the  existence  of 
an  element  with  any  particular  properties  was  regarded  as  an 
isolated  and  accidental  fact  in  nature  ;  but  the  periodic  arrange- 
ment presents  the  idea  that  it  is  necessary  that  elements  of 
given  atomic  mass  shall  have  certain  definite  properties. 

330.  "  Gaps  "  in  the  Periodic  Arrangement.  —  When 
Mendelejeff  first  drew  up  his  table  of  the  elements,  he 
found  that  in  several  cases  neighboring  heterologous  ele- 
ments did  not  fall  into  place,  i.  e.,  into  the  vertical  rows 
containing  the  other  members  of  their  natural  families. 
Thus,  the  element  zinc  (65)  was  followed  by  arsenic 
(75).  The  interval  is  large,  viz.,  ten  units.  Now,  zirc 
belongs  in  the  same  natural  family  with  magnesium  ;  ai;d 
if  arsenic  follows  zinc,  arsenic  must  belong  to  the  family 
of  boron  and  aluminum.  Hence  the  second  and  fourth 
horizontal  rows  would  be  as  follows :  — 

2.  Ka  (23)  Mg  (24)  Al  (27)  Si  (28)  P  (31)  S  (32)  Cl  (35.5) 
4.  Cu  (03)  Zn  (65)  As  (75)  Se  (79)  Br  (80),  etc. 

This  arrangement  is  manifestly  impossible  ;  for  by  all  its 
properties  bromine  belongs  with  chlorine,  selenium  with 
sulphur,  and  arsenic  with  phosphorus. 

The  arrangement  of  the  second  and  fourth  rows  should 
be,  — 

2.  Na  (23)  Mg  (24)  Al  (27)  Si  (28)  P  (31)  S  (32)  Cl  (35.5) 
4.  Cu  (63)  Zn  (65)  As  (75)  Se  (79)  Br  (80). 


PREDICTION  OF    UNKNOWN  ELEMENTS.        311 

There  were,  therefore,  two  gaps  in  the  fourth  row ;  a 
member  of  the  aluminum  family  and  a  member  of  the 
silicon,  family  were  wanting. 

331.  Prediction  of  Unknown  Elements.  — Reasoning 
from  the  assumption  that  the  properties  of  an  element 
are  determined  by  its  position  in  the  periodic  grouping, 
Mendelejeff  drew  up  a  table  stating  the  properties  of 
the  unknown  elements  that  ought  to  exist  to  fill  the  gaps 
in  the  fourth  row.  The  supposed  element  of  the  alu- 
minum family  he  called  ek-aluminum,  and  that  of  the 
silicon  family,  eka-silicon.  Another  gap  existed  in  the 
third  row,  between  calcium  (40)*and  titanium  (48).  To 
this  element  Mendelejeff  gave  the  provisional  name  eka- 
boron. 

Mendelejeff  and  Meyer's  tables  were  published  in 
1871.  Four  years  afterward  an  element  having  the 
properties  of  ek-aluminum  was  discovered  in  France,  and 
named  Gallium.  Its  atomic  mass  was  found  to  be  TO, 
as  predicted. 

In  1880,  Nilson  and  Cleve  discovered  in  a  Scandina- 
vian mineral  a  new  element  which  had  the  properties  of 
eka-boron.  They  called  it  Scandium. 

In  1886,  Clemens  Winkler  discovered  the  element 
which  he  called  Germanium.  Upon  comparing  the 
properties  of  this  element  (its  atomic  mass  is  72)  with 
those  foretold  for  eka-silicon,  he  decided  that  the  new 
element  could  be  nothing  else  than  the  eka-silicon  pre- 
dicted by  Mendelejeff  fifteen  years  before. 


312  THE  PEEIODIC  SYSTEM. 

332.  The  Table  of  the  Elements — The  periodic  system 

of  the  elements  is  given  on  the  following  page.  The  student  will 
notice  that  there  are  eight  instead  of  seven  groups.  This  is  due 
to  the  fact  that  there  are  three  groups  of  three  elements  each 
that  do  not  fit  into  the  table  of  seven.  The  first  of  these 
groups  consists  of  the  elements  iron,  cobalt,  and  nickel,  whose 
atomic  masses  place  them  between  manganese  (55)  and  copper 
(03).  These  three  groups  of  elements  constitute  Group  VIII. 

The  valence  of  the  elements  of  Group  IV  in  their  highest 
hydrogen  compounds  is  4;  from  Group  IV  to  Group  VII  the 
valence  toward  hydrogen  decreases. 

The  valence  of  the  elements  in  their  highest  oxygen  com- 
pounds increases  from  Group  I  to  Group  VII. 

M  is  the  general  symbol  of   all  the  elements. 

333.  Conclusion.  —  Besides  resulting  in  the  prediction 
of  the  properties  of  undiscovered  elements,  the  periodic 
classification  has  led   to   a  more   careful   study  of  the 
atomic  masses  of  known  elements. 

The  classification  has  many  faults,  but  it  is  full  of 
suggestions,  and  shows  a  striking  relationship  between 
the  elements.  Because  of  this  relationship  chemists  are 
tending  to  the  belief  that  all  the  elements  are  modifica- 
tions of  some  yet  simpler  forms  or  form  of  matter.  As 
to  the  character  of  this  fundamental  substance,  nothing 
is  known  at  present.  A  recent  irregularity  in  the  peri- 
odic arrangement  results  from  the  fact  that  the  atomic 
mass  of  tellurium  has  been  placed  at  about  127,  or  slightly 
higher  than  iodine,  126.9.  This  is  not  startling,  how- 
ever, for  the  periodic  recurrence  of  properties  is  only 
approximate,  at  the  best. 


THE  PERIODIC   TABLE. 


313 


1—  I 

s 

1 

oc 
1—  i 
CO 
CO 

CT5 
i—  i 
0 
GO 

cr 

GO 

0 

05 
CO 
05 

CO 
CO 

P 
to 

CO 

c 

—  ' 

M 

o 

o 

cj 

w 

CfQ 

s? 

2 

SP 

P 

O 

g 

w 

K 

Q 

i—  i  Ed 

to 

0 

k^ 
CO 

i—  i 

t—  * 

GO 

05 

o 

to 

CO 

o 

h-i  o 
cj 

0 

to 

O5 

^ 

H 

cr 

r 

i—  i 

p 

*j 

O 

0 

fe 

w 

g 

ng> 

to 

0 

i—  i 

CO 

M 

CO 
GO 

i—  ' 

GO 
CO 

0 

£ 

to 

(—1 

h-  ^ 

P 

'M§ 

p- 

cr 

? 

P° 

N 

0 

H 

03 

0 

R 

K 

MO 

to 

CO 

to 

to 

o 

i 

t—  * 
i—1 

CO 

P 

05 

to 

GO 

to 

00 

1—  I 

to 

o 

te 

*W 

^d 

w- 

^J 

8 

CO 

^ 

^ 

O 

P 

o^ 

cr 

ac 

^  ^ 

V 

^-\ 

^ 

s 

W 

to 

o 

GO 

1—  J 
GO 
CO 

to 
o 

cp 

2 

CO 
1—  I 

1—  I 

b 

Cn 

w 

'  3 

cj 

Si 

H 

CO 

K 

CO 

0 

Cfc 

0 

g 

s 

^ 

to 

CO 
CO 

1—  I 
00 

*•> 

to 

o 

CO 

to 

CO 

to 

1—  ' 

05 

0 

W 

^§ 

p' 

W 

g 

9 

*J 

^ 

^0 

H-i 

O 

to 

05 
CO 

1 

^ 

GO 
O 

P 
en 

CO 
Cn 

en 

1—  I 
CO 

M 

O 

-a 

w 

P§ 

2 

2 

N3 

1—  I 

""g  CO 

?$ 

^en 

<^o 

v» 
1—  I      L_J 

!-*  HH 

M<SO 

K  o 

CO  ^ 

0?  ^ 

S  0 

HH  Cj 

CO 

JW 

I—1 

0 

* 

314 


THE  PERIODIC   SYSTEM. 


334.  The  Argon  Family ._  The  discovery  of  argon  and 
the  elements  related  to  it,  viz.,  helium,  neon,  krypton,  and  xenon, 
has  led  to  the  introduction  of  a  new  family  into  Chemistry,  — 
the  Argon  family  (cf.  §  115).  Banisay  suggests  (November, 
1901)  that  the  group  may  be  placed  in  a  vertical  row  at  the  left 
of  the  alkali  group  in  the  table.  These  two  vertical  rows  would 
then  look  something  like  this  :  — 


He  4 

Li  7 

Ne  20 

Na  23 

Ar  40 

K39 

Cu  63 

Kr  82 

Rb  85.4 

Ag  108 

Xe  128 

Cs  133 

The  fact  that  argon  (40)  comes  before  potassium  (39)  need 
cause  no  more  anxiety  than  the  fact  that  tellurium  (127.5)  pre- 
cedes iodine  (126.9).  The  properties  of  the  elements  are  only 
approximately,  not  absolutely,  periodic  functions  of  the  atomic 
masses. 

The  elements  of  the  argon  group  are  all  monatomic ;  and 
their  valence  is,  apparently,  o. 


CHAPTER   XXI. 
SILICON   AND   BORON. 

A.    Silicon. 

335.  Occurrence  of  Silicon.  —  As  was  stated  in  §  6, 
silicon  is,  next  to  oxygen,  the  most  abundant  constituent 
of  the  earth's  crust.     It  is  not  found  free,  but  in    the 
form  of  its  oxide,  SiO2,  and  in  silicates,  i.  e.,  the  salts  of 
silicic   acid.     Silicon  dioxide   (silica)   and   the  silicates 
make  up  sand,  clay,  and  almost  all  the  crystalline  rocks 
of  the  earth's  crust. 

Silica  is  taken  up  by  plants.  The  straw  and  luisks  of  the 
grains  contain  it.  The  equisetum  is  called  k-  scouring-rush," 
from  the  large  amount  of  silica  present  in  it.  Silica  is  found, 
also,  in  the  skin,  nails,  hair,  etc.,  of  animals. 

Certain  microscopic  plants,  the  diatoms,  have  skeletons  of 
silica ;  and  these  have  accumulated  in  large  deposits  in  some 
places. 

336.  Preparation.  —  Silicon,   like   carbon,   exists   in 
several  allotropic  forms.     The  amorphous  variety  may  be 
obtained  by  heating  a  mixture  of  powdered  quartz  and 
powdered  magnesium. 

2  Mg  +  SiO2 »  2  MgO  +  Si. 

Amorphous  silicon  is  a  brown  powder  which  burns,  when 

515 


316  SILICON  AND  BORON. 

heated  in  the  air,  to  silicon  dioxide.     It  is  attacked  by  hydro- 
fluoric acid,  but  not  by  other  acids. 

Si  +  4  HF1  -  »  SiFl4  +  2  H2. 


Silicon  is  obtained  crystalline  by  heating  sodium  fluo- 
silieate,  Na2SiFl6,  with  sodium  and  zinc. 

Na2SiFl6  +  4  Na  -  »  6  NaFl  +  Si. 

The  silicon  dissolves  in  the  melted  zinc,  and  separates  out  in 
crystals  as  the  zinc  cools.  When  the  zinc  solidifies  it  encloses 
the  silicon.  The  zinc  is  then  removed  by  treating  the  mass 
with  hydrochloric  acid;  and  silicon  remains. 

The  crystalline  form  of  silicon  does  not  burn  in  air 
or  oxygen,  nor  does  it  dissolve  in  acids.  It  dissolves  in 
hot  sodium  hydroxide  solution,  giving  sodium  silicate  and 
hydrogen.  The  simplest  equation  is,  — 

Si  +  4  NaOH  -  »  Na4SiO4  +  2  H^ 

337.  Silicon  Compounds.  —  The  general  structure  oi 
silicon  compounds  is  like  that  of  the  corresponding  car- 
bon compounds.  Hydrogen  silicide,  SiH4,  corresponds  to 
marsh  gas,  CH4  ;  silicon  tetrachloride,  SiCl4,  to  carbon  tet- 
rachloride, CC14;  silicon  dioxide,  SiO2,  to  carbon  dioxide, 
CO2  ;  silicic  chloroform,  SiHCl3,  to  chloroform,  CHC13. 

Hydrogen  silicide  is  a  colorless  gas.  It  may  be  ob- 
tained, mixed  with  hydrogen,  by  treating  magnesium  sili- 
cide with  dilute  hydrochloric  acid. 

Mg2Si  +  4  HOI  -  »  SiH4  +  2  MgCl2. 


SILICON  COMPOUNDS.  317 

As  ordinarily  made,  it  ignites  spontaneously  in  the  air. 
SiH4  +  2  02 »  Si02  +  2  H2O. 

Magnesium  silicide  is  prepared  by  heating  powdered  quartz 
with  an  excess  of  magnesium  powder. 

4  Mg  +  Si02 »  Mg2  Si  +  2  MgO. 

If  too  little  magnesium  is  used,  amorphous  silicon  results 
(c/.  §  336). 

Silicon  tetrachloride  is  a  colorless  liquid,  boiling  at 
59°  C.  It  is  formed  from  silicon  and  chlorine.  Water 
decomposes  it,  giving  silicic  acid  and  hydrochloric  acid. 

SiCl4  +  4  H20 »  Si(OH)4  +  4  HC1. 

(Of.  the  action  of  phosphorus  trichloride  with  water, 
§  303.) 

Silicon  tetrafluoride  is  a  colorless  gas,  formed  when 
silicon  dioxide  is  treated  with  hydrofluoric  acid. 

Si02  +  4  HF1 »  SiFl4  +  2  H20. 

Silicon  tetrafluoride  is  decomposed  by  water  as  the  tetrachlor- 
ide is. 

SiFl4  +  4  H20 >  Si(OH)4  +  4  HF1. 

Instead  of  being  set  free,  however,  the  hydrofluoric  acid  unites 
with  some  of  the  silicon  tetrafluoride,  forming  fluosilicic  acid, 
H28iFl6.  The  name  "  fluosilicic  acid "  means  silicic  acid, 
H2SiO3,  with  its  oxygen  replaced  by  fluorine,  —  three  bivalent 
oxygen  atoms  by  six  univalent  fluorine  atoms  (c/.  §  314,  sulph- 
arsenites,  and  §  194,  thio sulphates).  Many  fluosilicates  are 
known;  potassium  fluosilicate,  K2SiFl6,  is  one  of  the  few  diffi- 
cultly soluble  potassium  salts. 


318  SILICON  AND  BORON. 

338.  Silicon  Carbide.  —  Silicon  carbide  (SiC)  or  car- 
borundum is  among  the  three  hardest  substances  known, 
the  others  being  boron  carbide  and  the  diamond.     It  is 
made  by  heating  a  mixture  of  powdered  quartz,  coke, 
saw-dust,  and  common  salt  in  the  electric  furnace. 

Carborundum  is  not  attacked  by  acids  nor  by  solutions  of 
alkalies.  It  burns  only  with  great  difficulty.  Because  of  its 
hardness,  powdered  carborundum  is  used  as  a  polishing  and 
cutting  agent. 

339.  Silicon  Dioxide.  —  Silicon  dioxide  is  found  as 
sand,  quartz,  etc.  (c/.  §  335).     The  pure  substance  may 
be  prepared  by  heating  silicic  acids,  H4SiO4,  H2SiO3,  etc. 

H4Si04  =  Si02  +  2  H20. 
H2Si03  =  Si02  +  H20. 

Silicon  dioxide  is  thus  silicic  anhydride. 

When  any  form  of  silicon  dioxide  is  fused  with  sodium  or 
potassium  hydroxide  or  carbonate,  sodium  or  potassium  silicate 
results. 

2  NaOH  -f  SiO2 »  Na2SiO3  -f  H2O. 

Ka2CO3  +  SiO2 »  Na2SiO3  +  CO2. 

Calcium  carbonate  acts  in  the  same  way. 

When  silica  is  fused  with  the  oxide  of  a  metal,  a  silicate  is 
also  formed. 

CaO  -f  SiO2  =  CaSiO3. 

Acids  do  not  act  upon  silica  (except  hydrofluoric  acid 
as  in  §  337). 

340.  Silicic  Acid.  — When  a  soluble  silicate  is  treated 
with  hydrochloric  acid,  a  bulky  mass,  like  gelatine,  is 


SILICATES.  319 

precipitated.  This  probably  consists  of  normal  silicic 
acid,  H4SiO4.  When  the  gelatinous  mass  is  dried  at 
the  ordinary  temperature,  it  loses  water  and  becomes  a 
non-crystalline  powder.  This  is  probably  ordinary  silicic 
acid,  H2SiO3.  When  the  powder  is  heated  to  a  high 
temperature  it  loses  water,  forming  silica,  SiO2.  The 
equations  are  given  in  §  339. 

Polysilicic  Acids.  —  Besides  the  normal  and  ordinary  forms 
of  silicic  acid,  many  other  forms  are  possible  ;  these  are  known 
as  polysilicic  acids.  They  are  derived  from  normal  silicic  acid 
by  the  loss  of  different  proportions  of  water.  A  general  for- 
mula for  them  all  would  be,  — 

o;Si(OH)4-yHsO. 
Thus,  if  x  =  2  and  y  =  1,  \ve  have,  — 

2H4Si04  —  H20  =  H6Sis07- 

Among  the  varieties  of  amorphous  silica  found  in  nature  are 
agate,  chalcedony,  opal,  cornelian.,  flint,  amethyst,  etc.  These 
all  contain  water,  and  may  therefore  be  looked  upon  as  forms 
of  the  polysilicic  acids.  The  colors  of  these  substances  are 
usually  due  to  traces  of  impurities. 

341.  Silicates.  —  The  silicates  are  salts  of  silicic  acid. 
The  mineral  chrysolite  is  magnesium  silicate,  Mg2SiO4. 
Serpentine  is  Mg3Si2O7.  These  are  salts  of  the  acids 
H4SiO4  and  H6Si2O7,  respectively.  Potassium,  sodium, 
and  calcium  silicates  are  derived  from  the  acid  H2SiO3. 

Potassium  silicate  is  known  as  "water  glass  "  ;  it  is  u£ed  to 
make  cements  and  artificial  stone.  Kaolin  is  practically  pure 
aluminum  silicate,  Al2(SiO3)3.  It  fuses  only  at  a  very  high 


320  SILICON  AND  BORON. 

temperature.  It  is  used  for  making  china  and  crockery  ware, 
fire-bricks,  etc.  Clay,  which  is  impure  aluminum  silicate,  melts 
lower  than  kaolin  ;  it  is  used  for  making  pottery,  bricks,  etc. 
The  red  color  of  baked  clay  is  due  to  ferric  silicate,  Fe2(SiO3)3. 

342.  Glass.  —  Glass  is  a  mixture  of  certain  silicates, 
generally  of  sodium  or  potassium  silicate  with  calcium 
or  lead  silicate. 

The  silicates  of  calcium,  lead,  etc.,  crystallize  when  they  sol- 
idify ;  a  glass  made  from  them  would  break  into  fragments  on 
cooling.  The  silicates  of  sodium  and  potassium,  however,  not 
only  do  not  crystallize  themselves,  but  even  prevent  the  other 
silicates  from  crystallizing. 

Ordinary,  soft  glass,  such  as  is  used  for  window 
panes  and  bottles,  is  essentially  a  mixture  of  the  silicates 
of  calcium  and  sodium;  it  may  be  made  by  melting  to- 
gether silica,  calcium  carbonate,  and  sodium  carbonate,  in 
the  proper  proportions. 

Hard  glass  is  a  mixture  of  the  silicates  of  calcium 
and  potassium.  It  is  used  for  making  chemical  appara- 
tus, lamp  globes,  etc. 

Flint  glass,  such  as  is  used  in  making  optical  instru- 
ments, etc.,  is  a  mixture  of  potassium  and  lead  silicates. 

Enameled  or  *  *  milky ' '  glass  is  made  by  adding 
cryolite  (cf.  §  266)  to  ordinary  glass. 

"  Granite  ironware  "  or  "porcelain-lined"  ware  con- 
sists of  iron  covered  with  an  easily  fusible  glass,  called 
enamel. 

Color  is  imparted  to  glass  by  the  addition  of  small  amounts 


PREPARATION   OF  BORON.  321 

of  other  substances.  Thus,  a  cobalt  compound  colors  glass 
blue  ;  a  cuprous  compound,  red  ;  a  chromic  compound,  green. 

The  etching  of  designs  on  glass  is  done  with  hydrofluoric 
acid,  as  described  in  §  268.  In  certain  kinds  of  etching  a  blast 
of  sand  is  used. 

Pressed  glassware  is  made  in  molds  ;  in  cut  glassware  the 
designs  are  ground  or  polished  by  means  of  emery,  carborun- 
dum, or  sandstone  wheels. 

All  articles  of  glass,  to  be  permanent,  must  be  an- 
nealed. Annealing  consists  in  allowing  the  hot  object 
to  cool  regularly,  so  that  its  molecules  may  assume  per- 
manent positions  with  reference  to  one  another.  Unan- 
nealed  glass  may  fly  to  pieces  at  the  slightest  jar. 

B.  Boron. 

343.  Occurrence  of  Boron.  —  The  element   boron  is 
the  first  member  of  the  aluminum  group   of  elements 
(<?/".  page  313) ;  yet  in  its  free  state  it  closely  resembles 
silicon  and  carbon.     It  occurs  in  nature  chiefly  as  boric 
acid  (H3BO3)  and  as  borax  (Na2B4O7). 

Boric  acid  is  found  chiefly  in  Tuscany,  where  it  issues,  mixed 
with  steam,  from  the  earth.  Borax  is  found  in  Nevada  and 
California  in  dried  borax  lakes.  Boracite,  (Mg3B8O15)2.  MgCl2, 
occurs  at  Stassfurt,  in  Germany. 

344.  Preparation.  —  Boron  exists   in   two  allotropic 
forms,  —  the  amorphous  and  the  crystalline  ;  it  is  difficult 
to  prepare  either  form  in  a  pure  state. 

Crystalline  boron  is  made  by  heating  boron  trioxide,  B2O3, 
with  aluminum.  The  aluminum  reduces  the  oxide  ;  and  the 


322  SILICON  AND  BORON. 

liberated  boron  dissolves  in  the  excess  of  aluminum.  When 
the  cooled  mass  is  treated  with  hydrochloric  acid,  the  aluminum 
dissolves,  leaving  crystals  of  boron. 

345.  Properties.  —  Crystalline  boron  has  a  specific 
gravity  of  2.6,  and  resembles  the  diamond.  The  amorph- 
ous form  is  brown.  It  burns  when  heated  in  the  air, 
giving  boron  trioxide,  B2O3.  Nitric  acid  converts  it  into 
boric  acid.  The  crystalline  form  is  more  difficult  to 
ignite  than  the  amorphous. 

Boron  dissolves  in  melted  sodium  hydroxide,  giving  soc^jum 
borate.  When  it  is  heated  in  nitrogen  or  ammonia  it  gives 
boron  nitride,  BN.  This  is  a  white  powder  which  is  decom- 
posed by  steam,  giving  boric  acid  and  ammonia. 

BN  +  3  H2O >  B(OH)3  ,-f  NH3. 

With  chlorine,  boron  forms  boron  trichloride,  BC13,  a  colorless 
liquid.  This  is  decomposed  by  water,  giving  boric  acid  and 
hydrochloric  acid  (c/.  §§  303  and  337). 

BC13  +  3  H20 >  B(OH)3  +  3 


346.  Boric  Acid.  —  Boric  acid  is  made  by  adding 
concentrated  hydrochloric  or  nitric  acid  to  a  hot  solution 
of  borax.  It  is  a  wliite,  crystalline  solid.  Its  aqueous 
solution  has  a  faintly  acid  reaction  with  litmus,  but 
colors  turmeric  paper  brown,  just  as  sodium  hydroxide 
does..  Boric  acid  is  in  some  respects  like  aluminum  liy- 
droxide,  which  is  an  acid  or  a  base,  according  to  circum- 
stances. 


EXERCISES.  323 

Boric  acid  is  volatile  with  steam,  as  was  indicated  in  §  343t 
Its  solution  in  ethyl  alcohol  burns  with  a  bright  green  flame. 
AVhen  boric  acid  is  heated  it  loses  water,  giving  finally  the  an- 
hydride, B2O3.  This  redissolves  readily  in  water,  giving  the 
acid. 

347.  Borax.  —  The  ordinary  borates  are  derived  not 
from  the  normal  acid,  H8BO8,  but  from  tetraboric  acid, 
H2B4O7.     This  is  related  to  the  normal  acid  just  as  the 
poly  silicic  acids  are  to  normal  silicic  acid  (</.  §  340). 

4  H3B03  —  5  H20  =  H2B407. 

Borax  is  sodium  tetraborate,  Na2B4O7.  It  comes 
into  the  market  as  crystals  which  are  either  Na9B4O7. 
10  H2O  or  Na2B4O7.5  HaO. 

When  heated,  borax  loses  its  crystal-water  and  swells  to  a 
porous  mass  ;  on  stronger  heating,  this  melts  to  a  clear  glass. 
Molten  borax  dissolves  the  oxides  of  metals,  and  with  some  of 
them  gives  characteristic  colors  ;  hence  its  use  in  the  laboratory 
to  form  the  "  borax  bead,"  and  in  the  arts  to  clean  the  surfaces 
of  metals  before  soldering  and  welding,  and  before  coating 
metals  with  enamel  (of.  §  342). 

Borax  is  used,  also,  to  prevent  the  decomposition  of  organic 
substances,  in  medicine,  in  the  manufacture  of  hard-water 
soaps,  and  to  increase  the  gloss  of  starch. 

348.  Exercises. 

1.  Name  some  of  the  scouring  mixtures  which  consist  es- 
sentially of  silica. 

2.  Suggest  &  possible  reason  for  the  fact  that  silicon  dioxide 
is  a  hard,  infusible  solid,  while  the  corresponding  oxide  of  car- 
bon is  "raseous. 


324  SILICON  AND  BOB  ON. 

3.  Name  some  polishing,  or  abrasive,  agents  besides  silica 
and  carborundum.     How  are  diamonds  polished  ? 

4.  From  some  of  the  facts  named  in  this  chapter  suggest  a 
way  in  which  carbon  dioxide  might  be  liberated  from  calcium 
carbonate  in  the  earth's  crust. 

5.  What  class  of  substances  must  be  present  in  underground 
waters  so  that  silica  may  be  held  in  solution  ?     What  substance 
is  needed  to  hold  calcium  carbonate  in  solution  ? 

6.  From  what  polysilicic  acid  would  a  silicate  of  the  formula 
CaSi2O5  be  derived  ?     Show  the  relation  of  this  acid  to  normal 
silicic  acid. 

7.  What  are  the  proportions,  by  volume,  of  factors  and  prod- 
ucts in  the  equation  representing  the  combustion  of  hydrogen 
silicide  (c/.  §  337)  ? 


CHAPTER   XXII. 
DISSOCIATION   AND   MASS   ACTION. 

349.  Dissociation    by    Heat.  —  We    have    already 
learned  of  several  cases  of  dissociation  by  heat,  e.  #.,  the 
decomposition  of  steam  (ef.  §  45)  and  of  hydriodic  acid 
(<?f.  §  276).     The  essential  things  in  a  dissociation  are 
the  constant  decomposition  of  complex  molecules  into  sim- 
pler ones,   and   the  constant  recombination  of   the  sim- 
pler molecules,  until  a  definite  equilibrium  is  established. 

The  dissociation  of  iodine  (c/.  §  274)  is  recognized  by  the 
fact  that  while  below  600°  C.  the  molecular  mass  of  iodine 
vapor  is  254,  above  this  temperature  the  molecular  mass  dimin- 
ishes, until  at  about  1500°  C.  it  becomes  practically  127.  This 
diminution  of  the  molecular  mass  means  that  the  number  of 
molecules  has  increased.  !„  molecules  have  become  monatomic 
(I)- 

350.  Dissociation  in  Solution  ;  lonization.  —  Disso- 
ciation takes  place  not  only  at  elevated  temperatures, 
but,  in  many  cases,  in  solution.     It  is  recognized  just  as 
heat-dissociation  ia,  viz.,  by  an  increase  in  the  number  of 
molecules.     As  we  have  already  learned  (cf.  §§  244  and 
245),  the  molecular  mass  of  a  soluble  substance  can  be 
obtained  from  the  osmotic  pressure,  and  from  the  freez- 
ing-point and  the  boiling-point  of  the  solution.     When, 
however,  we  come  to  apply  these  methods  to  a  substance 

325 


326  DISSOCIATION  AND  MASS  ACTION. 

like  sodium  chloride,  we.  notice  that  the  molecular  mass 
is  too  small,  and  in  dilute  solution  only  about  one-half 
of  what  Ave  should  expect. 

To  illustrate  :  An  aqueous  solution  of  cane-sugar  (C12H22OU) 
containing  342  grams  of  sugar  to  the  liter  (342  is  the  molecular 
mass  of  cane-sugar)  freezes  at  about  — 1.8°  C.;  so  does  a  solu- 
tion of  grape-sugar  (C6H12O6),  containing  180  grams  of  sugar 
to  the  liter.  But  the  freezing-point  of  a  solution  of  sodium 
chloride  containing  58.5  grams  of  salt  to  the  liter  approaches 
— 3.6°  C.  ;  the  depression  is  thus  twice  as  great  as  we  should  ex- 
pect. This  can  be  due  only  to  the  presence  of  twice  as  many 
molecules  as  we  should  expect ;  the  molecules  of  sodium  chlor- 
ide must  have  dissociated  in  aqueous  solution. 

Many  other  substances,  e.  g.,  hydrochloric  acid,  nitric 
acid,  potassium  and  sodium  hydroxides,  potassium  chloride, 
sodium  nitrate,  etc.,  are  like  sodium  chloride.  The  de- 
pression of  the  freezing-point,  the  elevation  of  the  boiling- 
point,  and  the  osmotic  pressure  of  aqueous  solutions  of 
these  substances  all  show  that  the  number  of  molecules 
present  in  solution  is  greater  than  we  should  expect — 
in  dilute  solution  practically  twice  as  great.  The  mole- 
cules of  sodium  sulphate,  sulphuric  acid,  etc.,  dissociate 
into  three  particles  in  dilute  solution. 

The  electric  conductivity  of  solutions  supports  the 
theory  of  dissociation ;  for  the  electric  current  can  pass 
readily  only  through  solutions  containing  substances 
whose  molecules  dissociate  in  solution.  Pure  water 
has  a  conductivity  that  is  almost  nil ;  the  conductivity 
of  sugar  solution  is  also  very  small ;  but  aqueous  solu- 


EXPLANATION   OF  NEUTRALIZATION.          327 

tions  of  strong  adds,  strong  bases,  and  the  salts  formed 
from  them  conduct  readily.  The  substances  just  named 
are,  therefore,  called  electrolytes,  and  dissociation  in  so- 
lution is  called  electrolytic  dissociation.  The  particles 
into  which  an  electrolyte  dissociates  are  called  ions ; 
hence  the  dissociation  is  called,  also,  ionization. 

Composition  of  a  Solution.  —  A  solution  of  sodium  chloride  in 
water  is,  then,  not  merely  a  mixture  of  the  molecules  of  these 
two  substances.  In  the  first  place,  most  of  the  sodium  chloride 
molecules  are  dissociated  into  sodium  and  chlorine  ions.  Since 
recombination  goes  on  at  the  same  time,  the  behavior  of  sodium 
chloride  in  water  is  shown  by  the  equilibrium  equation  (c/.  § 

276),— 

I          

NaCl  ~ »  Na  +  Cl. 

But  the  water,  too,  is  dissociated  slightly, — 

+        - 
IL,Q- »H-fOH. 

Hence  combination  may  take  place  between  sodium  and  hy- 
droxyl  ions  and  between  hydrogen  and  chlorine  ions  to  form 
uridissociated  sodium  hydroxide  and  hydrochloric,  acid,  respec- 
tively. We  have,  therefore,  present  in  an  aqueous  solution  of 
sodium  chloride  eight  distinct  things  :  the  ions  hydrogen,  hy- 
droxyl,  sodium,  and  chlorine,  and  the  undissociated  substances 
water,  sodium  chloride,  sodium  hydroxide,  and  hydrochloric 
acid. 

351 .  Explanation  of  Neutralization.  —  A  very  simple 
explanation  of  the  properties  of  acids,  bases,  and  salts, 
and  of  neutralization  follows  from  the  ionization  theory. 
All  acids  in  aqueous  solution  affect  litmus  in  the  same 


328  DISSOCIATION  AND   MASS  ACTION. 

way  because  they  all  give  hydrogen  ions.  The  turning 
of  litmus  red  is  simply  an  indication  of  the  presence  of 
these  ions.  Similarly,  bases  are  substances  whose  solu- 
tions contain  hydroxyl  (OH)  ions. 

Neutralization  is  thus  essentially  the  union  of  hydrogen 
and  hydroxyl  ions  to  form  undissociated  molecules  of  water. 

This  new  definition  of  neutralization  is  supported  by  the  fact 
that  the  amount  of  heat  evolved  when  the  several  strong 
acids  are  neutralized  by  a  given  base  is  approximately  the 
same  if  equal  numbers  of  molecules  of  the  acids  are  taken. 
Thus,  63  grams  of  nitric  acid,  36.5  grams  of  hydrochloric  acid, 
and  49  grams  of  sulphuric  acid  all  liberate  practically  the  same 
quantity  of  heat  when  neutralized  by  sodium  hydroxide.  Sul- 
phuric acid  has  the  molecular  mass  98  ;  but  its  molecule  gives 
two  hydrogen  ions  ;  hence  49  grams  of  sulphuric  acid  are  equiva- 
lent to  63  grams  of  nitric  acid  and  to  36.5  grams  of  hydrochloric 
acid. 

The  ionic  equation  for  the  neutralization  of  sodium 
hydroxide  by  hydrochloric  acid  is,  — 


fa  +  OH  +  H  +  Cl »  Na  +  Cl  +  H2O. 

When  we  evaporate  the  water  we  get  union  of  the  so- 
dium and  chlorine  ions ;  and  sodium  chloride  is  formed. 

352.  "  Test "  Reactions  are  Ionic. —  What  we  call 
tests  for  classes  of  substances,  e.  a.,  the  action  of  silver 
nitrate  solution  with  chlorides,  and  of  barium  chloride 
solution  with  sulphates,  are  really  tests  for  free  ions. 
Thus,  when  we  add  silver  nitrate  solution  to  a  solution 
of  sodium  chloride,  the  equation  is,  — 


ELECTROLYSIS.  329 

Ag  +  N03  +  Na  +  Cl  -  tAgCl  +  Ka  +  KO3. 


The  silver  chloride,  being  practically  insoluble,  is  almost 
all  removed  as  a  precipitate. 

If,  however,  we  add  silver  nitrate  solution  to  a  dilute 
solution  of  sodium  chlorate,  NaClO3,  we  do  not  get  a 
precipitate  of  silver  chloride,  because  no  chlorine  ions  are 
present.  The  chlorine  of  sodium  chlorate  is  part  of  the 
ion  ClOy  Silver  chlorate,  which  might  be  formed,  is 
not  precipitated  because  soluble.  Hence  no  result  is 
apparent. 

+  +        -  +     •  -        + 

Ag  +  N03  +  Na  +  C108  -  >  Ag  +  N03  +  Na  +  CIO,. 

Another  illustration  :  Ferrous  chloride,  FeCl2,  gives  with 
ammonium  sulphide,  (NH4)2S,  a  black  precipitate  of  ferrous  sul- 
phide, FeS  ;  but  ammonium  sulphide  does  not  precipitate  the 
iron  of  potassium  ferrocyanide  solution,  K4Fe(GN")6,  because 
iron  ions  are  not  present. 

K4Fe(CN)6  -  »  4  K  +  Fe(CK)6. 

353.  Electrolysis.  —  What  we  call  electrolysis  is  not 
the  tearing  apart  of  molecules  by  the  electric  current, 
but  the  carrying  of  electric  charges  by  the  ions  from  one 
electrode  to  the  other.  When  we  electrolyze  dilute  sul- 
phuric acid,  as  in  §  37,  we  have,  before  the  introduction 
of  the  electrodes  into  the  acid,  ionization  of  most  of  the 
sulphuric  acid  molecules. 

H8SO4  -  ^ 


330  DISSOCIATION  AND  MASS  ACTION'. 

• 

The  current  passes  because  the  hydrogen  ions  carry 
-J-  charges  from  the  -j~  to  the  —  electrode,  and  the  sul- 
phuryl  (SO4)  ions  —  charges  from  the  —  to  the  -f-  elec- 
trode. The  bivalent  >S'04  ion  can  carry  as  great  a  charge 
as  two  univalent  hydrogen  ions. 

+ 
The  charged  ions  (H  and  SO4)  are  not  atoms ;  neither  have 

they  all  the  properties  of  free  molecules.  When  they  reach 
their  respective  electrodes,  however,  they  take  on  their  ordinary 
properties.  The  hydrogen  ions  give  up  their  .-f-  charges  at  the  — 
electrode,  and  become  neutral  atoms.  These  unite  to  form  hv- 
drogen  molecules.  The  sulphuryl  ion  gives  up  its  —  charge 
at  the  -j-  electrode,  and  becomes  the  radical  SO4.  This,  being 
unstable,  acts  upon  water,  giving  sulphuric  acid  and  atomic 
oxygen. 

SO4  -f  H2O » II2SO4  +  — O— . 

The  oxygen  atoms  unite  to  give  molecules  of  oxygen  (O2)  or 
of  ozone  (O3 ;  cf.  §  287). 

354.  Hydrolysis.  —  While  strong  acids,  like  hydro- 
chloric acid,  etc.,  are  much  dissociated  in  solution,  weak 
acids,  like  carbonic  acid,  hydrocyanic  acid,  etc.,  are  dis- 
sociated only  slightly.  Consequently,  when  we  dissolve 
the  salt  formed  from  a  weak  acid  and  a  strong  base,  e.  g., 
sodium  carbonate,  we  get  the  reaction  of  the  base,  i.  e., 
of  OH  ions. 

That  this  must  be  so  is  seen  from  the  equation, 

+        -          +  +  - 

2  Xa  +  C03  +  2  II  +  2  Oil »  II2CO3  +  2  Na  +  2  OH. 

The  carbonic  acid  is  only  slightly  dissociated  ;  hence  relatively 
few  hydrogen  ions  remain  in  the  solution,  while  the  hydroxyl 


MASS  ACTION.  331 

ions  are  present  in  excess.     Therefore  sodium  carbonate  solu- 
tion reacts  alkaline. 

The  dissociation  of  the  salts  formed  by  weak  bases  and 
strong  acids  is  similar ;  only  in  this  case  we  get  the  reac- 
tion of  the  acid,  i.  e.,  of  hydrogen  ions. 

Thus,  when  ferric  chloride  is  dissolved  in  water,  the  ionic 
equation  is, — • 

Fe  +  3  Cl  +  3  H  +  3  OH »  Fe(OH)8  -f  3  II  +  3  Cl. 

Here  the  ferric  hydroxide,  Fe(OH)3,  being  a  weak  base  and  little 
dissociated,  removes  many  OH  ions  from  the  "sphere  of  ac- 
tion "  ;  hence  the  H  ions  determine  the  reaction  of  the  solution. 
Ferric  chloride  solution  reacts  acid.  Morever,  its  rusty  color 
shows  that  much  undissociated  ferric  hydroxide  (iron  rust)  is 
present  in  it. 

The  illustrations  just  given  are  cases  of  hydrolysis, 
i.  e.,  of  "  decomposition  by  water." 

355.  Mass  Action. — The  recombination  of  dissoci- 
ated particles,  whether  in  gaseous  form  or  in  solution,  is 
influenced  by  the  frequency  with  which  the  particles 
meet.  If,  therefore,  we  wish  to  stop  or  to  diminish  the 
dissociation  of  a  substance  AB,  we  simply  see  to  it  that 
a  large  excess  of  one  of  the  dissociated  particles,  e.  g., 
A,  is  present.  The  active  mass  of  A,  that  is,  the  mass 
of  it  that  can  combine  with  B,  is  thus  increased. 

The  effects  due  to  an  excess  of  dissociated  particles 
are  cases  of  Mass  Action. 

To  illustrate :  When  we  attempt   to  get  the  molecular 


332  DISSOCIATION  AND  MASS  ACTION. 

mass  of  phosphorus  pentacJdoride,  PC16,  by  vapor  density 
methods,  we  find  that  the  molecular  mass  is  too  low,  owing  to 
the  dissociation  of  some  of  the  molecules  into  phosphorus 
trichloride  and  chlorine. 

PC15 »  PC13  +  C12. 

The  dissociation  can  be  prevented  if  the  vapor  density 
determination  is  carried  out  in  an  atmosphere  of  phosphorus 
trichloride ;  because  the  chlorine  molecules  then  meet  mole- 
cules of  phosphorus  trichloride  so  frequently  that  practically 
no  chlorine  molecules  remain  free. 

Similar  effects  occur  in  solution.  Thus,  if  \ve  wish  to  pre- 
cipitate the  sulphuric  acid,  i.  e.,  the  S04  ions,  contained  in  a 
solution,  we  use  barium  chloride  solution.  If  we  use  exactly 
the  theoretical  amount  of  barium  chloride,  however,  much  ba- 
rium sulphate  remains  in  solution  (as  J3a  and  SO4  ions).  The 
reaction  is  only  partly  represented  by  the  equation, 

Ba  +  2  Cl  +  2  II  +  S04 »  BaSO4  +  2  II  +  2  Cl, 

for  the  reverse  operation  also  goes  on. 

i         

BaSO4 »  Ba  -f  SO4. 

Whatever  will  cause  the  $04  ions  to  meet  more  Ba  ions  in  a 
given  time  will  increase  the  precipitation  of  the  $04  ions  as 
barium  sulphate.  We  obtain  this  result  by  adding  a  large  excess 
of  barium  chloride  solution. 


356.  Exercises. 

i.  What  substances  are  present  in  a  solution  of  potassium 
nitrate  in  water  ?     Why  is  the  reaction  neutral  ? 


EXERCISES.  333 

2.  Make  a  definition  of  an  acid  in  terms  of  the  ionization 
theory.     Of  a  salt. 

3.  The  reaction  of  a  solution  of  potassium  cyanide,  KCN, 
is  alkaline  ;  so  is  that  of  a  solution  of  disodium  hydrogen  phos- 
phate, Na2HPO4.     Explain. 

4.  Write  the  ionic  equation  for  the  neutralization  of  hydro- 
bromic  acid,  HBr,  by  potassium  hydroxide.     Explain. 

5.  Cold,  concentrated  sulphuric  acid  does  not  act  upon  zinc, 
but  the  dilute  acid  does."     Explain. 

6.  Explain,  in  terms  of  the  ionic  theory,  the  action  of  hydro- 
yen  sulphide,  H28,  upon  cupric  sulphate  solution  (c/.  §  182). 

7.  To  precipitate  all  the  manganese  of  a  manganese  sulphate 
solution  as  sulphide  (  cf.  §  182),  we  use  a  decided  excess  of  am- 
monium sulphide,  (KH4)2  S.     Why  ? 


CHAPTER   XXIII. 
METALS. 

357.  Metals   and   Non-Metals.  —  In   our    previous 
work   we    have    studied   chiefly   non-metallic    elements. 
There  is,  however,  110  sharp  distinction  between  metals 
and  non-metals,  but  rather  a  gradual  change  from  one 
class   to   the  other   (cf.   §  327).     But  just   as    oxygen, 
chlorine,  and  sulphur  have  a  distinct  character,  which  no 
one   would  mistake  for   that  of  a  metal,  so  there   are 
certain  elements  having  a  typical  metallic  nature. 

Metals  are  usually  opaque,  and  their  polished  surfaces 
are  good  reflectors  of  light ;  hence  the  metallic  luster. 
They  are  good  conductors  of  heat  and  electricity.  With 
oxygen  and  hydrogen  the  metals  form  bases;  and  by 
replacing  the  hydrogen  of  acids,  i.  e.,  ionic  hydrogen, 
they  form  salts. 

Some  elements,  e.  g.,  antimony,  are  both  acid-forming  and 
base-forming  ;  they  are  often  called  metalloids. 

358.  Occurrence  of   Metals. --The  solid   elements 
and  compounds  found  in  nature  are  called  minerals. 
The    minerals   and  mixtures    of   minerals    from  which 
metals  are  obtained  are  called  ores.     Some  metals,  e.  g., 
gold  and   copper,  occur  free,   that  is,   uncombined  with 
other  elements  ;  but  most  metals  are  found  as  oxides  or 

334 


PROPEBTIES   OF   THE  METALS.  335 

sulphides.  Some  metals  are  found  as  carbonates,  hydrox- 
ides, etc. 

359.  Extraction  of  Metals  from  Their  Ores.  —  If 

metals  occur  free  they  may  be  separated  by  mechanical 
means  from  minerals  mixed  with  them.  An  illustration 
is  the  crushing  of  an  ore  in  a  stamp-mill  and  the  wash- 
ing away  of  the  lighter  materials.  Copper  and  gold  are 
extracted  in  this  way,  although  chemical  methods  are 
employed  with  inferior  ores  of  these  metals. 

The  most  common  method  of  extracting  metals  is  to 
reduce  the  oxide  with  carbon  (charcoal  or  coke).  This 
is  the  case  with  iron.  If  the  ores  used  are  not  oxides, 
they  are  usually  converted  into  oxides  by  "  roasting " 
(cf.  §  178).  Sulphides,  hydroxides,  and  carbonates  may 
thus  be  changed  into  oxides. 

Another  method  of  reducing  an  oxide  is  to  heat  it  in  a  stream 
of  hydrogen.  The  oxygen  and  hydrogen  unite  and  escape  as 
steam,  while  the  metal  is  left.  This  is  a  good  laboratory 
method;  but  it  is  too  expensive  for  commercial  use. 

Chlorides  are  sometimes  reduced  by  heating  them  with 
sodium.  Aluminum  was  formerly  obtained  in  this  wa}r  from 
its  chloride. 

Several  of  the  metals,  e.  g.,  aluminum,  are  obtained 
by  the  action  of  the  electric  current  upon  some  of  their 
compounds. 

360.  Properties  of  the  Metals. — Besides  the  gen- 
eral  properties  already  mentioned,  the  metals  possess 


336  METALS. 

other  properties  in  varying  degrees.  Thus,  some  metals, 
e.  g.,  sodium  and  lead,  are  soft ;  while  others,  e.  g.,  chrom- 
ium and  manganese,  are  hard.  Sodium  and  lithium  are 
light  enough  to  float  on  water,  while  gold  is  19.3  and 
platinum  21.5  times  as  heavy  as  water. 

Again,  some  metals  evolve  much  energy  in  uniting  with 
oxygen,  while  others  form  very  unstable  oxides.  Usually  the 
lighter  metals,  such  as  sodium  and  potassium,  are  very  active 
chemically,  and  form  strong  bases  ;  while  the  heavy  metals , 
such  as  lead  and  gold,  are  much  less  active. 


CHAPTER   XXIV. 


THE   ALKALI  METALS. 

361.  General  Properties.  — The  metals,  like  the  non- 
metals,  are  generally  studied  in  groups  or  natural  fami- 
lies based  upon  similarity  of  properties.  The  Alkali 
group  consists  of  the  jive  metals  named  below  and  -the 
radical  ammonium,  NH^  the  compounds  of  which  re- 
semble those  of  sodium  arid  potassium  (cf.  §  148). 
These  metals  are  called  "  alkali "  metals  because  the 
two  most  important  members  of  the  group  are  contained 
in  the  alkalies,  i.  e.,  in  sodium  and  potassium  hydroxides. 


ELEMENT. 

SYMBOL. 

ATOMIC 

MASS. 

SPECIFIC 
GRAVITY. 

MELTING 
POINT. 

Lithium. 

Li 

7 

0.59 

180°  C. 

Sodium. 

Ka 

23 

0,97 

95.  6°  C. 

Potassium. 

K 

39 

0.87 

62.5°  C. 

Rubidium. 

Rb 

85.4 

1.52 

38.5°  C. 

Caesium. 

Cs 

133 

1.85 

26.  5°  C. 

337 


338  THE  ALKALI  METALS. 

All  of  these  metals  have  a  silvery  white  luster  and 
are  easily  cut.  In  air  they  become  coated  with  a  layer 
of  the  oxide  and  the  hydroxide ;  if  carbon  dioxide  is 
present,  these  pass  into  the  corresponding  carbonates. 
The  alkali  metals  burn  when  heated  in  air,  and  decom- 
pose water  at  ordinary  temperatures ;  therefore  none  of 
them  is  found  free  in  nature.  The  salts  of  these  metals 
are  practically  all  soluble  in  water. 

The  properties  of  the  alkali  metals  change  in  the  order  of 
the  atomic  masses,  e.  (/.,  the  higher  the  atomic  mass  the  lower 
the  melting-point.  The  chemical  activity  and  the  electro-posi- 
tive character  increase  from  lithium  to  ccesium.  Caesium  is  the 
most  electro-positive  element  known. 

362.  Lithium.  —  Lithium    is  widely  distributed   in 
nature,  but  no  mineral  known  contains  a  large  propor- 
tion of   it.     It  is  found   in  minute  quantities  in  most 
mineral    waters,    in    many   plants,    and    in    the    blood. 
Lithium  is  the  lightest  of  the  metals.     Its  salts  color 
the  Bunsen  flame  crimson. 

363.  Sodium.. —  Sodium   occurs   widely   distributed 
and  in  large  quantities,  especially  as   sodium  chloride, 
NaCl.     This  exists  as.  rock-salt  and  sea-salt,  and  in  many 
mineral  springs.     Sodium  silicate  is  found  in  many  rocks ; 
the  nitrate  is  Chili  saltpeter.     Large  deposits  of  the  sul- 
phate and  carbonate  exist.     The  ashes  of  plants  grow- 
ing in  or  near  the  sea  contain  sodium  carbonate,  and 
were  formerly  the  source  of  many  sodium  compounds. 


PREPARATION  AND  PROPERTIES  OF  SODIUM.    339 

364.  Preparation  and  Properties  of  Sodium.  —  At 

present,  sodium  is  prepared  by  the  electrolysis  of  sodium 
hydroxide,  or  of  sodium  chloride.  Strontium  chloride 
or  potassium  chloride  is  added  to  the  sodium  chloride  to 
lower  the  melting-point. 

Formerly  the  metal  was  made  by  heating  a  mixture  of  the 
carbonate  and  charcoal  in  the  absence  of  air. 

Na2CO3-j-  2  C  -  »  2  !Na  -f  3  CO. 

The  sodium  which  distilled  off  was  first  condensed  to  a  liquid, 
and  was  then  collected  under  petroleum. 

Sodium  is  a  white,  soft  metal  that  can  be  moulded  be- 
tween the  fingers,  and  can  be  readily  pressed  into  wire. 
Although  so  soft  at  ordinary  temperatures,  it  is  quite 
hard  at  —  20°  C.  It  decomposes  water,  producing 
sodium  hydroxide  and  hydrogen  (cf.  §  46)  ;  this  method 
is  used  to  prepare  pure  sodium  hydroxide. 

Sodium  unites  readily  with  oxygen  (cf.  §  289),  producing  a 
mixture  of  the  monoxide  (Na2O)  and  the  peroxide  (Na2O2). 
The  properties  of  the  pure  monoxide  are  not  known.  The  per- 
oxide has  recently  come  into  general  use  as  an  oxidizing  and 
bleaching  agent  and  as  a  source  of  free  oxygen.  It  is  decom- 
posed by  water,  giving,  finally,  sodium  hydroxide  and  oxygen. 


(1)  Ka2O2  +  2  H20  -  »  2  NaOH  +  H2O2. 

(2)  2H202  =  2 


An    alloy    of   sodium   and    mercury,    called    sodium 
amalgam,  is  a  useful  reducing  agent  ;  it  is  simply  diluted 


340  THE  ALKALI  METALS. 

sodium.     An  amalgam  is  an  alloy  of  which  mercury  is 
one  constituent. 

365.  Sodium    Hydroxide.  —  Sodium    hydroxide,    or 
caustic  soda,   is  formed  when   sodium   or  either  of  its 
oxides  acts  upon  water.     Two  of  the  commercial  methods 
of  preparing  the  hydroxide  are  as  follows  :  - 

(1)  Boiling  a  solution  of  sodium  carbonate  with  slaked 
lime  (calcium  hydroxide). 

Na2CO3  -|-  Ca(OH)2 >  2  NaOH  +  CaCOs. 

The  calcium  carbonate,  being  insoluble,  is  precipi- 
tated. The  sodium  hydroxide  solution  is  drawn  off 
and  evaporated. 

(2)  Electrolysis  of  a  concentrated  solution  of  sodium 
chloride. 

2  ^aCl  -f  2  H20 »  2  NaOII  +  1I2  +  C12. 

The  sodium  hydroxide  and  hydrogen  collect  at  the 
—  electrode. 

Sodium  hydroxide  is  a  white,  deliquescent  solid,  very 
soluble  in  water.  Both  the  solid  and  its  solution  absorb 
carbon  dioxide  readily.  The  solution  has  a  soapy  feel- 
ing, and  turns  red  litmus  blue. 

366.  Soap.  —  Sodium  hydroxide  is  one  of  the  strong- 
est and  most  useful  bases.     When  fats  are  boiled  with 
it  they  are  saponified,  i.  e.,  converted  into  soap.     The  fats 
are  chiefly  glyceryl  salts  of  organic  acids.  They  are  decom- 
posed by  sodium  hydroxide  into  glycerine  and  the  organic 
acid.     The  sodium  salt  of  the  organic  acid  is  soap. 


SODIUM  CARBONATE.  341 

C  1  -lie. 


C8HS^0  — CO  —  c"lC    +    HONa 


O  —  CO  —  C17II35 
glyceryl  stearate  sodium  hydroxide 

.Oil 

C3H5  —  OH  -f  3  NaO.  CO.  C17H8B. 
XOH 

glycerine       sodium  stearate 
(soap) 

The  sodium  stearate  formed  is  "salted  out"  of  solution  by 
adding  sodium  chloride. 

When  soap  is  dissolved  in  water  it  is  partly  hydrolyzed  (  cf. 
§  354  )  into  sodium  hydroxide  and  the  organic  acid  ;  the  sodium 
hydroxide  is  the  cleansing  agent.  "When  soap  is  put  into  hard 
water,  an  insoluble  scum  is  formed  ;  this  is  the  calcium  salt  of 
the  organic  acid. 

367.  Sodium  Carbonate.  —  Sodium  carbonate,  or  soda, 
is  one  of  the  most  important  chemicals  manufactured. 
The  principal  methods  of  making  it  are  the  Solvay,  or 
Ammonia,  Process  and  the  Le  Blanc  Process.  The 
second  of  these  is  the  more  interesting  historically,  be- 
cause it  was  devised  by  Le  Blanc  for  the  French  gov- 
ernment during  the  Revolution,  when  the  supply  was 
cut  off;  but  the  ammonia  process  is  so  much  cheaper 
that  fully  three-fourths  of  the  soda  used  is  now  made 
in  this  way.  Both  processes  begin  with  common  salt. 

The  Solvay  Process  consists  essentially  in  treating  sodium 
chloride  with  ammonium  hydrogen  carbonate,  NH4HCO3. 

KaCl  +  NH4HC03 »  NH4C1+  NaHCOs. 

This  reaction  takes  place  because  sodium  hydrogen  carbonate 


342  THE  ALKALI  METALS. 

(sodium  bicarbonate)  is  not  very  soluble  in  water  and  is,  there- 
fore, readily  precipitated  when  ammonium  hydrogen  carbonate 
is  added  to  concentrated  salt  solution.  The  solution  of 
ammonium  hydrogen  carbonate  is  formed  by  passing  carbon 
dioxide  under  pressure  into  a  saturated  solution  of  ammonia 


H2CO3  -  >  NH4HC03  +  H2O. 
Gentle  heating  converts  the  bicarbonate  into  carbonate. 
2  NaHCOg  =  NajCO3  +  H2O  -f  CO2. 

The  carbon  dioxide  thus  set  free  is  used  again  ;  and  nearly 
all  the  ammonia  is  recovered  by  heating  the  brine,  from  which 
the  bicarbonate  has  crystallized,  with  slaked  lime  (cf.  §  142). 

The  Le  Blanc  Process  consists  of  essentially  three  opera- 
tions :  — 

(1)  The  conversion  of  common  salt  into  sodium  sulphate  (cf. 
§  92)  ; 

2  NaCl  +  H2S04  -  »  Na2S04  +  2  HC1. 

(2)  The  reduction  of  sodium  sulphate  to  sodium  sulphide; 

Na2S04  +  40  -  »  Ka2S  -f  4  CO. 

(3)  The  conversion  of  sodium  sulphide  into  the  carbonate; 
2S  +  CaC03  -  >  Ka2CO3  +  CaS. 


The  second  and  third  operations  take  place  together,  sodium 
sulphate  being  mixed  with  limestone  and  coal-dust  and  the 
mixture  heated.  The  sodium  carbonate  cannot  readily  be  sepa- 
rated from  calcium  sulphide  because  both  are  soluble.  If  lime- 
stone is  present  in  excess,  however,  some  of  it  is  dissociated 
into  quicklime  (CaO)  and  carbon  dioxide  ;  the  quicklime  and 
the  calcium  sulphide  form  an  insoluble  compound. 

Soda  comes  into  the  market  as  calcined  soda,  or  soda- 
ash,  containing  no  crystal-water,  and  as  crystallized  soda, 


SODIUM  CHLORIDE.  343 

or  sal  sodae,  Na2CO3.  10  H2O.  Soda  is  used  in  great 
quantities  in  the  manufacture  of  glass  (c/.  §  342)  and 
of  sodium  hydroxide  (§  365). 

368.  Sodium  Bicarbonate.  —  Sodium  bicarbonate  is 
prepared  by  treating  the  normal  salt  with  carbonic  acid 
(</.  §  210).     The    equation    for  its   decomposition  by 
heat  is  given  in  §  367.     It  is  used  in  large  amounts  for 
baking  powders,  effervescing  mixtures,  "  soda-water,"  and 
medicine. 

369.  Sodium  Phosphate.  —  The  common  phosphate 
of  sodium  is   disodium   hydrogen  phosphate,  Na2HPO4. 
12  H2O  ;  it  is  formed  by  adding  sodium  carbonate  to 
phosphoric  acid  until  the  solution  is  slightly  alkaline. 

Phosphoric  acid  is  tribasic  (c/.  §  304)  and  its  insoluble  salts 
are  usually  normal,  e.  g.,  Li3PO4,  Ca3(PO4)2,  Ag3PO4  ;  but  its 
normal  sodium  salt  (Na3PO4)  is  so  readily  hydrolyzed  (c/.  §  354) 
that  it  absorbs  carbonic  acid  from  the  air. 

2  Ka3P04  -f  H2CO3 »  2  Na2HPO4  -f  Xa2CO3. 

370.  Sodium   Chloride.  —  Sodium  chloride,  or  com- 
mon salt,  occurs  widely  distributed.     It  makes  up  about 
3^>  of  sea-water,  and  is  found  in  large  deposits  in  Gali- 
cia  (Austria),    Germany,   England,  the   United  States, 
etc.     In  some  places  salt  is  mined  as  rock-salt,  while  in 
others  the  mixture  of  salt  and  earth  is  treated  with  water, 
the  resulting  brine  being  pumped  to   the  surface  and 
then  evaporated.     At  Manistee,  Mich.,  the  brine  is  con- 
centrated and  the    salt    continuously    separated    in   a 
"  vacuum  "  boiler. 


344  THE  ALKALI  METALS. 

Sodium  chloride  crystallizes  in  colorless,  transparent 
which  decrepitate  (cf.  §  50)  when  heated.     It  is  only  a  little 
more  soluble  in  hot  than  in  cold  water  (cf.  §  59). 

Salt  is  necessary  to  the  life  of  man  and  other  animals,  the 
hydrochloric  acid  of  the  gastric  juice  being  derived  from  it 
(cf.  §  90).  It  is  used  in  enormous  quantities,  not  only  as  food, 
but  as  the  starting  material  in  the  preparation  of  most  com- 
pounds of  sodium  and  of  chlorine. 

371.  Sodium    Nitrate.  —  Sodium    nitrate,    or    Chile 
saltpeter,  occurs  in  enormous  quantities  in  the  Atacama 
Desert,  in  Chile  (jrf.  §163).     It  is  very  deliquescent, 
and  s,o  cannot  be  used  for  gunpowder,  etc.     It  is  con- 
verted into  potassium  nitrate,  which  is  not  deliquescent 
(<?/'.  §  164),  and  into  nitric  acid  (c/.   §  155).     Sodium 
nitrate  is  also  a  constituent  of  artificial  fertilizers.     The 
crude  salt  is  an  important  source  of  iodine  (cf.  §  273). 

372.  Other    Sodium    Salts.  —  Sodium    sulphate,    or 
"  Glauber's  Salt,"  Na2SO4.  10  H2O,  is  made  from  com- 
mon salt  and  sulphuric  acid  (</.  §  92).     It  is  made,  also, 
from  magnesium  sulphate  and  salt. 

MgSO4  -f  2  NaCl  — -»  Na2SO4  -f  MgCl2. 

This  salt  is  very  efflorescent  (cf.  §  53)  ;  its  principal  use  is  as 
a  substitute  for  soda  in  the  manufacture  of  glass. 

Sodium  thiosulphate,  Na2S2O3.  5  H2O,  has  been  de- 
scribed in  §  194 ;  and  sodium  tetraborate  (borax)  in 
§347. 

Sodium  salts  color  the  Bunsen  flame  bright  yellow. 


POTASSIUM  HYDltOXIDE.  345 

373.  Potassium.  — The  element  potassium  is  a  con- 
stituent of  many  minerals ;  among  these  are  feldspar,  a 
double    silicate    of   potassium    and    aluminum;    sylvite, 
practically  pure  potassium  chloride;  potash  alum;  and 
saltpeter. 

Potassium  is  present  in  all  soils  ;  it  doubtless  comes  from 
the  disintegration  of  feldspar  and  other  rocks.  Plants  take 
up  potassium  compounds;  hence  the  element  is  found  (as  the 
carbonate,  potash)  in  plant  ashes.  When  the  ashes  are  ex- 
tracted with  water,  potassium  carbonate  dissolves. 

Potassium  is  made  by  the  electrolysis  of  potassium 
cyanide,  hydroxide,  or  chloride.  It  is  a  soft  metal  like 
sodium,  but  has  a  slightly  bluish  luster.  It  decomposes 
water,  giving  the  hydroxide  and  hydrogen.  The  energy 
evolved  in  the  reaction  is  so  great  that  the  escaping  gas 
is  set  on  fire. 

The  vapors  of  potassium  and  of  all  potassium  salts 
color  the  flame  violet.  Potassium,  like  sodium,  must  be 
kept  under  kerosene  or  ligroin  to  protect  it  from  moist 
air. 

374.  Potassium  Hydroxide.  —  Potassium  hydroxide, 
or  caustic  potash,  is  made  by  the  action  of  the  electric 
current  upon  a  concentrated  solution  of  potassium  chlor- 
ide (>/.  §  365). 

2  KC1  -f  2  H2O >  2  KOH  +  H2  -f  C12. 

It  may  also  be  made  from  potassium  carbonate  and 
"  milk  of  lime,"  i.  e.,  calcium  hydroxide. 


346  THE  ALKALI  METALS. 

-  K2CO3  +  Ca(OH)2 »  2  KOH  -f  CaCOz. 

Potassium  hydroxide  is  a  white,  deliquescent  solid 
It  is  a  powerful  base.  It  absorbs  carbon  dioxide  from 
the  air,  forming  potassium  carbonate. 

375.  Potassium  Carbonate,  or  Potash.  —  Much  potas- 
sium carbonate  is  prepared  by  the  Le  Blanc  and  Solvay  processes 
from  potassium  chloride  (c/.  §  367).     The  crude  substance  is 
obtained  from  wood-ashes.     Anhydrous  potassium  carbonate  is 
a  powerful  dehydrating  agent  (c/.  §  53). 

Potash  is  used  chiefly  in  making  the  hydrjxide  and  hard 
glass  (c/.  §  342). 

376.  Potassium  Nitrate.  —  Potassium  nitrate  (called, 
also,  saltpeter  and    nitre)  has    already    been    described 
(<?/.  §§  163  to  165).     Most  of  that  in  use  is  made  from 
sodium  nitrate  and  potassium  chloride,  both  of  which  are 
found    in   large   deposits.     The    equation    is   given   in 
§  164. 

377.  Potassium  Chlorate.  —  Potassium  chlorate  re- 
sults when  chlorine  is  passed  into  a  hot,  concentrated 
solution   of   potassium  hydroxide  until  the  solution  is 
saturated  (cf.  §  281). 

6  KOH  -f  3  C12 >  5  KC1  +  KC1O3  -f  3  H2O. 

A  cheaper  way  is  to  pass  the  chlorine  into  hot  "  milk  of 
lime  "  ;  calcium  chlorate,  Ca(ClO3)2,  is  formed.  This  with  po- 
tassium chloride  gives  potassium  chlorate  and  calcium  chloride. 

Ca(ClO8)2  -f  2  KC1 »  2  KC109  +  CaCl2. 


POTASSIUM  BROMIDE  AND  POTASSIUM  IODIDE.    347 

Potassium  chlorate,  being  much  less  soluble  in  cold  water  than 
the  other  substances,  crystallizes  out,  leaving  the  others  in 
solution. 

Like  the  nitrate,  potassium  chlorate  is  valuable  chiefly 
as  an  oxidizing  agent.  It  is  used  in  preparing  oxygen, 
explosive  mixtures,  e.  g.,  smokeless  powder,  matches,  and 
fireworks.  It  is  sold  by  druggists  as  "  potash  "  for  sore 
throats.  * 

378.  Potassium  Bromide  (KBr)  and  Potassium 
Iodide  (KI).  —  Potassium  bromide  and  potassium  iodide 
may  be  prepared  by  the  action  of  bromine  and  iodine,  re- 
spectively, upon  potassium  hydroxide.  The  equations 
are  analogous  to  the  one  for  the  action  of  chlorine 
upon  caustic  potash. 

6  KOH  -f  6  Br  -  »  5  KBr  +  KBrO3  +  3  H2O. 
6  KOH  -f  6  I  -  »  5  KT  -f  KTO3  +  3  H2O. 

By  evaporating  the  solution  containing  bromide  and 
bromate,  or  iodide  and  iodate,  to  dryness,  and  then  heat- 
ing the  residue  sufficiently,  we  can  decompose  the  brom- 
ate and  the  iodate  just  as  we  can  the  chlorate  (cf. 

§19). 

2  KBrO3  =  2  KBr  -f  3  O2. 


Potassium  bromide  and  iodide  are  made,  also,  by  treating  the 
bromide  and  the  iodide  of  iron  with  potassium  carbonate. 

Fe3Br8  -f  4  K2CO3  -  >  Fe3O4  -f-  8  KBr  -j-  4  CO2. 
Fe3I8  %  4  K2C03  -  >  Fe304  -f  8  KI  -f  4  CO2. 


348  THE  ALKALI  METALti. 

The  iron  compounds  are  formed  by  adding  bromine  and 
iodine,  respectively,  to  moist  iron  turnings. 

379.  Ammonium.  —  -  The    formation    of    ammonium 
salts   by  neutralizing  acids   Avith  ammonium  hydroxide 
has    already    been    described   (cf.  §  148).     The   radical 
NH±  has  not  been  isolated  because  it  decomposes  into 
ammonia  and  hydrogen.      One  of  the  many  arguments 
for  its  existence  is  the  formation  of  ammonium  amalgam 
by  the  action  of  a  strong  solution  of  ammonium  chloride 
upon  sodium  amalgam  (cf.  §  364). 

Ka,  Hg  -f  NTI4C1 »  !N"H4,  Hg  -f  NaCl. 

sodium  ammonium 

amalgam  amalgam 

Ammonium  amalgam  is  a  bulky,  metallic  mass  re- 
sembling sodium  amalgam.  It  decomposes  readily  into 
ammonia,  hydrogen,  and  mercury. 

380.  Exercises. 

1.  Why   does   the    electrolysis   of    an   aqueous  solution   of 
sodium  chloride   give  sodium  hydroxide  and   hydrogen   at  the 

-  electrode  (cf.  §  365)  ? 

2.  Write  the  equations  showing  how  potassium  carbonate  is 
made  by  the  Solvay  process  (cf.  §  375). 

3.  All  four  substances  given  in  the  equation  in  §  372  are 
soluble.     Explain  under  what  conditions  the  reaction  can  take 
place. 

4.  Describe  the  preparation  of  ammonium  chloride,  nitrate, 
and  sulphate,  giving  equations.     How  does  ammonium  nitrate 
behave  when  heated  ?    Ammonium  chloride  ? 


EXERCISES.  349 

5.  What  is  the  source  of  the  ammonia  of  commerce? 

6.  How  do   you  explain  the  fact  that  sodium  bicarbonate 
solution  reacts  alkaline? 

7.  Which  of  the  following  gases  would  you  dry  with  solid 
caustic  potash  :  ammonia,  hydrogen  sulphide,  carbon  monoxide, 
carbon  dioxide,  oxygen? 


CHAPTER   XXV. 
THE  ALKALINE-EARTH  METALS. 

381.  The  Group.  —  The  "alkaline-earth"  metals  are 
so   called   because   they  form   the   transition  from  the 
alkalies  to  the  "  earth  "  metals,  such  as  aluminum.     In 
this  chapter  we  shall  consider  glucinum  (or  beryllium), 
magnesium,  calcium,  strontium,  and  barium.     The  most 
important  is  calcium. 

382.  Calcium  (Atomic  Mass,  40) Calcium  does  not 

occur  free  ;  but  its  compounds  are  found  in  large  quan- 
tities.     The    most   abundant  is  the   carbonate,  CaCO3  ; 
this  occurs  as  limestone,  marble,  chalk,  calc-spar,  and  coral. 
The  sulphate,  CaSO4,  the  phosphate,  Ca3(PO4)2,  and  the 
fluoride,  CaFl2,  are  also  important  miners. 

The  metal  is  obtained  by  heating  calcium  Made  with  carbon 
in  a  stream  of  Irydrogen  in  an  electric  furnace.  It  is  a  silvery 
solid.  It  decomposes  water  like  sodium  and  potassium. 

383.  Calcium    Oxide    (CaO).  —  Calcium    oxide    is 
familiar  to  all  as  lime,  or  quicklime.     It  is  made  by  heat- 
ing calcium  carbonate  in  large  furnaces  called  lime-kilns. 

CaCO3  =  CaO  -f  CO2. 

Lime  is  a  white,  amorphous  solid,  fusible  only  at  the  tem- 
perature of  the  electric  furnace.  For  its  use  in  the  Drum- 

350 


CALCIUM  HYDROXIDE. 


mond,  or  lime,  light  see  §  11.  It  absorbs  water  and  carbon 
dioxide  from  the  air,  forming  air-slaked  lime,  which  consists 
Of  the  carbonate-  and  hydroxide.  Lime  is  thus  a  good  agent 
for  removing  water  and  carbon  dioxide  from  gases  (cf.  §  223). 

When  lime  is  treated  with  a  suitable  quantity  of 
water,  the  two  unite  to  form  a  soft,  dry  powder.  The 
operation  is  called  "  slaking  "  ;  and  the  product,  calcium 
hydroxide,  is  called  "  slaked  lime." 

CaO  +  H20  =  Ca(OH)2. 

So  much  heat  is  evolved  when  water  and  lime  unite, 
that  lime  improperly  protected  from  water  has  been  the 
cause  of  many  fires. 

384.  Calcium  Hydroxide  Ca(OH)2.  —  Calcium  hy- 
droxide is  not  very  soluble,  less  than  1%  parts  dis- 
solving in  1,000  of  water.  The  solution  is  lime-water; 
its  uses  have  been  considered  in  §§  99  and  209.  If 
more  of  the  hydroxide  is  present  then  the  water  will 
dissolve,  the  liquid  appears  milky  ("  milk  of  lime  "  ;  cf. 
§§  374  and  377). 

Uses.  —  Calcium  hydroxide  is  generally  prepared  from  the 
oxide  just  before  using.  It  is  used  to  prepare  ammonia  (cf.  § 
142),  the  hydroxides  of  sodium  and  potassium  (cf.  §§  365  and 
374),  bleaching  powder  (cf.  §§88  and  279),  chlorates  (cf.  §  377), 
glass,  mortar,  and  cements  (cf.  §  389);  to  purify  suqar  and  illumi- 
nating gas  (cf.  §  223)  :  to  remove  hair  from  hides  ;  to  extract 
metals  from  their  ores  :  as  a  disinfectant  and  white-wash,  and 
in  making  stearin  candles. 


352  THE  ALKALINE-EAETH  METALS. 

385.  Calcium  Chloride  (CaCl2) .  —  Calcium  chloride 
is  usually  made  from  calcium  carbonate  and  hydrochloric 
acid  (cf.  §  203).     The  anhydrous  substance   (made  by 
heating  at  200°  C.)  is  very  deliquescent,  and  dissolves 
in  water  with  evolution  of  heat.     It  is  used  as  a  drying 
agent  (<?f.  §  53). 

The  crystallized  chloride,  CaCl2.  6  H2O,  absorbs  heat  when 
dissolving  in  water  (cf.  §  57),  and  when  mixed  with  ice  or  snow 
may  produce  a  temperature  of  —  40°  C.  A  concentrated  solu- 
tion of  calcium  chloride  freezes  so  much  lower  than  one  of 
sodium  chloride  that  it  is  often  used  instead  of  brine  as  the 
cold  bath  in  ice  factories  (cf.  §  146). 

386.  Calcium     Carbonate    (CaC03).  —  As    already 
stated  (cf.  §  211)   calcium  carbonate   occurs    in  many 
forms  and  widely  distributed.     Limestone,  chalk,  ealcite, 
and  marble  are  found  in  large  masses.     Aragonite  crystal- 
lizes in  a  form  different  from  that  of  the  other  varieties, 
but  has  the  same  composition. 

Limestone,  the  most  abundant  form  of  calcium  carbonate,  is 
gray,  and  often  contains  small  crystals,  but  is  always  mixed 
with  clay  and  other  impurities.  Limestone  containing  much 
clay  is  marl.  Marl  is  used  in  making  cement.  Limestone  is 
used  as  a,  flux  (cf.  §  266)  in  smelting  iron,  as  a  source  of  lime, 
and  as  building-stone.  Chalk  is  used  for  making  lime.  Car- 
penters use  it  for  marking.  Marble  is  used  for  building 
purposes,  as  a  source  of  lime,  etc. 

387.  Calcium  Sulphate  (CaS04) .  —  Calcium  sulphate 
occurs  principally  as  gypsum,  CaSO4.  2  H2O.     Alabaster 
is  a  granular  form  of  gypsum. 


CALCIUM  PHOSPHATE.  353 

When  gypsum  is  heated  to  120°  to  130°  0.  it  loses 
about  three-fourths  of  its  crystal-  water,  and  forms  the 
white  powder  known  as  Plaster  of  Paris. 

2  CaSO4(II,0)a'=  (CaS04)2H2O  +  3  II2O. 
gypsum  plaster  of 

Paris 

When  this  powder  is  mixed  with  enough  water  to  form 
a  paste  it  expands,  and  hardens  to  a  mass  with  a  smooth 
surface.  Because  of  these  properties  plaster  of  Paris  is 
used  to  make  casts,  as  a  wall  finish,  and  as  a  cement. 
The  union  of  the  powder  with  water  produces  the  crys- 
talline compound. 


(CaSO4)2H2O  +  3  H2O  =  2  CaSO4(H2O)2. 

Calcium  sulphate  is  slightly  soluble  in  water  ;  water  contain- 
ing it  is  permanently  hard  (<-/.  §  43).  When  &  soluble  carbonate 
is  added  to  a  solution  of  calcium  sulphate,  calcium  carbonate  is 
precipitated. 


CaSO4  +  Ka2CO3  -  >  CaC09 
Hence  the  "  softening"  of  hard  water  by  means  of  soda. 

388.  Calcium  Phosphate,  Ca8(P04)2.  —  Normal  cal- 
cium phosphate  occurs  as  phosphorite;  combined  with 
calcium  chloride  or  fluoride  it  forms  apatite.  Important 
deposits  of  these  minerals  are  found  in  Florida  and 
South  Carolina.  Calcium  phosphate  is  the  chief  inor- 
ganic constituent  of  bones  (cf.  §  292).  Being  insoluble, 
normal  calcium  phosphate  is  converted  into  the  primary 


354  THE  ALKALINE-EARTH  METALS. 

phosphate,  Ca(H2PO4)2,  for  use  as  a  fertilizer  ((/. 
§  307).  The  primary  phosphate  is  commonly  known 
as  "  soluble  phosphate " ;  the  mixture  of  the  primary 
phosphate  and  calcium  sulphate  is  called  "super- 
phosphate." 

Phosphates  Necessary  for  Plants. — To  be  fertile,  soil  must 
contain  calcium  phosphate,  an  essential  plant  food.  When  the 
crops  are  removed,  part  of  the  phosphate  of  that  region  goes  with 
them  ;  hence,  phosphate  must  be  returned  to  the  soil  if  the 
land  is  to  yield  good  harvests.  If  the  crops  are  used  as  food 
for  animals,  part  of  the  phosphate  returns  to  the  soil  in 
manure  ;  if  not,  other  fertilizers  must  be  used.  Nature  usually 
keeps  a  soil  fertile  by  means  of  decaying  vegetation,  which 
forms  with  the  soil  "  vegetable  mould." 

Fertilizers.  — A  complete  fertilizer  supplies  potassium,  nitro- 
gen, and  phosphorus.  Most  fertilizers,  however,  contain  only 
one  or  two  of  these  essentials. 

Potassium  is  usually  returned  to  the  soil  as  the  sulphate  or 
carbonate  (wood-ashes  ;  cf.  §  378)  ;  sometimes  as  chloride. 

Nitrogen  is  frequently  Supplied  as  ammonium  salts,  or  as 
nitrates,  especially  sodium  nitrate  (cf.  §  371).  Nitrogen  is  also 
contained  in  guano. 

Phosphorus  is  contained  in  fertilizers  chiefly  as  "soluble 
phosphate,"  which  is  obtained  by  treating  phosphate  rocks  or 
bone-ash  with  sulphuric  acid. 

The"  dry  residue  left  after  the  waste  products  of  slaughter- 
houses, e.  </.,  tainted  meat,  bones,  hoofs,  etc.,  are  deprived  of 
fat,  oil,  and  gelatine,  makes  a  good  fertilizer.  The  fat  is  used 
as  soap-stock. 

389.  Mortar  and  Cement.  —  When  a  thick  paste  of 
slaked  lime  and  water  is  exposed  to  the  air  it  gradually 


OTHEE    CALCIUM  COMPOUNDS.  355 

"sets,"  i.  e.,  becomes  dry  and  hard.  The  chemical  ac- 
tion that  takes  place  consists  in  the  escape  of  water  and 
the  absorption  of  carbon  dioxide  from,  the  air. 

Ca(OH)2  +  H2CO3  (i.  e.,  H2O  -f  CO2) »  CaCO3  +  2  H2O. 

During  the  "  setting  "  the  mass  contracts.  In  the  mak- 
ing of  mortar  and  cements  this  contraction  is  overcome 
by  the  use  of  sand.  Sand  also  makes  the  mortar  more 
porous,  so  that  carbon  dioxide  can  penetrate  farther  and 
moisture  escape  more  easily.  Freshly  plastered  walls 
remain  mois,t  for  some  time  because  of  the  slow  libera- 
tion of  water  by  carbon  dioxide.  The  complete  change 
requires  a  long  time. 

When  limestone,  magnesium  carbonate,  and  clay  (aluminum 
silicate,  cf.  §  341)  are  heated  together,  there  is  formed  a  cement 
which  slakes  slowly  and  without  much  heat  evolution  ;  this 
cement  has  the  valuable  property  of  hardening  under  water. 
It  is  known  as  hydraulic  cement. 

Blast-furnace  slay  (cf.  §  430)  and  some  anhydrous  .silicates 
form  with  lime  ar  similar  cement. 

Portland  Cement  is  made  by  heating  to  a  very  high  tempera- 
ture an  artificial  mixture  of  calcium  carbonate  and  clay. 

390.  Other  Calcium  Compounds.  —  Bleaching  pow- 
der, CaOCl2,  has  already  been  described  in  §§88  and 
279.  It  was  formerly  supposed  to  be  a  mixture  of  cal- 
cium hypochlorite,  Ca(OCl)2,  and  calcium  chloride,  CaCl2. 
Such  a  mixture  would  have  the  same  quantitative  com- 
position as  bleaching  powder,  and  would  give  the  same 
ions  in  solution. 


356  THE  ALKALINE-EARTH  METALS. 

"When  a  concentrated  solution  of  bleaching  powder  is  boiled,  it 
gives  off  oxygen. 


O2. 


The  presence  of  small  amounts  of  the  oxides  or  hydroxides 
of  nickel  and  cobalt  makes  the  decomposition  more  easy. 

A  dilute  solution  of  bleaching  powder  gives,  when  boiled, 
the  chlorate  and  the  chloride  of  calcium. 


6  CaOCl2  =  Ca(ClO3)2  +  5  CaCl 


This  method  of  making  the  chlorates  was  considered  in  §  377. 

Calcium  silicate,  CaSiO3,  is  a  constituent  of  ordinary  glass 
(cf.  §  342).  It  may  be  made  by  adding  a  solution  of  sodium 
silicate  (water-glass)  to  a  solution  of  calcium  chloride,  or  by 
fusing  quartz  (SiO2)  with  calcium  carbonate. 

Calcium  carbide,  CaC2,  has  been  described  under  acetylene 
(§  221). 

Calcium  sulphide,  CaS,  is  a  white,  soluble  solid  made  by  re- 
ducing calcium  sulphate,  CaSO4.  It  is  a  by-product  in  the  Le 
Blanc  soda  process  (cf.  §  367).  After  commercial  calcium  sul- 
phide has  been  exposed  to  sunlight,  it  emits  a  phosphorescent 
glow  ;  hence  it  is  used  in  making  luminous  paints  for  match- 
boxes, clock-faces,  etc. 

391.  Barium  and  Strontium.  --  The  compounds  .of 
barium  and  .of  strontium  resemble  those  of  calcium;  but 
they  are  less  abundant  and  useful. 

The  monoxides  unite  with  water,  forming  the  hydrox- 
ides ;  these  are  string  bases,  like  calcium  hydroxide. 

Barium  peroxide,  BaO2,  is  formed  when  the  monoxide 
is  heated  to  dull  redness  in  air.  Its  use  as  a  source  of 
oxygen  was  given  in  §  21,  and  as  a  source  of  hydrogen 
peroxide,  in  §  289. 


MAGNESIUM  COMPOUNDS.  357 

Baiium  salts  impart  a  green  color  to  the  Bunsen  flame,  and 
are  poisonous.  The  nitrate,  Ba(NO3)2,  is  used  in  making  green 
lights,  fireworks,  etc.  The  chloride,  BaCl2.  2  II2O,  is  a  white, 
crystalline  solid.  "With  a  solution  of  a  sulphate  it  forms  the 
very  insoluble  barium  sulphate,  BaSO4,  and  thus  serves  as  a 
test  for  £04  ions  (cf.  §  189). 

Strontium  salts  impart  a  red  color  to  the  flame.  The 
nitrate,  Sr(NO3)2,  is  used  in  pyrotechny  for  producing 
red  lights.  The  hydroxide  is  used  in  refining  beet- 
sugar. 

392.  Magnesium.  —  Magnesium  occurs  as  the  chlor- 
ide, carbonate,  and  silicate.     The  metal  is  prepared  by 
the   electrolysis  of    the   double  chloride    of    magnesium 
and  potassium.     This  occurs  as  the  mineral  carnallite, 
MgCl2.  KC1.  6  H20. ' 

Magnesium  is  a  silvery,  white  metal  with  a  high  luster. 
It  is  light  (S.G.  1.74),  and  oxidizes  slowly  in  ordinary 
air.  When  heated  in  oxygen  it  burns  with  a  bright 
flame,  forming  magnesium  oxide,  MgO  (cf.  §  23)  ;  in  air 
it  burns  to  the  oxide  and  the  nitride,  Mg3N2  (cf.  §  115). 
The  metal  does  not  decompose  water  at  ordinary  tem- 
peratures; it  differs  in  this  respect  from  calcium,  stron- 
tium, and  bafium,  and  from  the  alkali  metals. 

Magnesium  powder  is  used  for  flash-lights  and  in  fire- 
works. 

393.  Magnesium    Compounds.  —  Magnesium     oxide 
(magnesia)  is  formed  Avhen  magnesium  burns  and  when 
the  carbonate  and  hydroxide  are  heated.     It  is  very  re- 


358  THE  ALKALINE-EARTH  METALS. 

fractory,  i.  e.,  hard  to   melt,  and    is   used  for  making 
crucibles,  cupels,  etc. 

Magnesium  chloride  forms  deliquescent  crystals 
(MgCl2.  6  H2O).  When  its  solution  is  evaporated  to 
dryness,  hydrochloric  acid  escapes.  The  residue  con- 
sists chiefly  of  the  oxide. 

MgCl2  +  H20  -  »  MgO  +  2  1IC1. 

Magnesium  sulphate,  MgSO4,  is  called  "  Epsom  salts." 
It  is  found  in  many  springs  and  in  the  mineral  Jcieserite, 
MgSO4.  H2O.  Kieserite  changes  in  contact  with  water 
to  the  salt  MgSO4.7  H2O,  which  is  the  .magnesium  sul- 
phate of  commerce. 

Magnesium  carbonate,  MgCO3,  occurs  as  the  mineral 
magnesite,  and,  combined  with  calcium  carbonate,  in 
dolomite. 

The  normal  carbonate  is  not  precipitated  by  alkaline  carbon- 
ates, owing  to  the  ease  with  which  it  is  hyclrolyzed  (cf.  §  354)  ; 
the  precipitate  consists  of  a  basic  carbonate.  At  the  same 
time  carbonic  acid  is  set  free. 


MgCl2  +  Na2CO3  -  »  MgCO3  -f  2  NaCl. 
MgC03  +  2  H20  -  »  Mg(OH)2  +  H2C03. 

The  basic  carbonate  is  a  mixture  of  the  carbonate  and  the 
hydroxide. 

Asbestos    is    a    mixture    of    calcium    and   magnesium 
silicates. 

394.  Glucinum  or  Beryllium  (Atomic  Mass,  9.08.)  - 
Glucinum    (Gl)    is   a  rare,  white   metal,   resembling  magnes- 


THE   GROUP  A   "NATURAL   FAMILY.' 


359 


him.     Its  compounds  have  a  sweetish  taste  ;  hence  the  name 
glucinum  (cf.  glucose,  glycerine,  etc.). 

395.  The  Group  a  "  Natural  Family. "  —  The  mem- 
bers of  the  calcium  group,  like  the  alkali  metals,  form 
a  natural  family  of  elements.  This  will  be  apparent 
from  a  comparison  of  some  of  the  properties  as  given 
in  the  following  table.  Magnesium  really  forms  with 
glucinum,  xzinc,  cadmium,  and  mercury  a  separate  divi- 
sion ;  but  many  of  its  properties  ally  it  with  the  calcium 
group. 


ELEMENT. 

Magnesium 

Calcium. 

Strontium. 

Barium. 

ATOMIC  MASS. 

24 

40 

88 

137 

SPECIFIC  GRAVITY. 

1.7 

1.6 

2.5 

3.6 

CARBONATE    DISSOCI- 
ATES ;  TEMPERATURE. 

300°  C. 

600°  C. 

1100°  C. 

1400°  C. 

GRAMS  OF  HYDROXIDE 
SOLUBLE  IN  A  LITER 
OF  WATER  AT  15°  C. 

0.009 

1.32 

18 

50 

HEAT  OF  FORMATION 
OF  CHLORIDE  ;  UNITS. 

151 

170 

185 

195 

360  THE  ALKALINE-EARTH  METALS. 

396.  Exercises. 

1.  The  temperature  at  which  strontium  and  barium  carbon- 
ates dissociate  being  very  high  (see  table),  suggest  how  to  get 
the  oxides  of  these  metals  more  easily  (c/.   §§  169  and  321). 
Write  a  possible  equation. 

2.  Why  do  carpenters  still  use  chalk  instead  of  crayon  for 
marking  ?     Crayon  consists  of  gypsum,  etc. 

3.  Magnesium  oxide  unites  with  water  much  less  readily  than 
calcium  oxide  ;  how,  probably,  would  strontium  oxide  compare 
with  calcium  oxide  ?     With  barium  oxide  ? 

4.  Write  and  explain  the  ionic  equation  for  the  action  of  a 
solution  of  ammonium  carbonate,  (NH4)2CO8,  upon  barium  chlor- 
ide solution. 

5.  How  many  grams   of   carbon   dioxide   can   be  obtained 
from  200  grams  of  pure  Iceland  spar  (cf.  §  211)?     How  many 
liters  at  20°  C.  and  740  mm.? 


CHAPTER   XXVI. 
ZINC,  CADMIUM,  AND   MERCURY. 

397.  The  Zinc  Group.  —  The  elements  zinc,  cadmium, 
and  mercury  are  members  of  the  second  group  of  the 
periodic  system.     They  are,  however,  much  more  closely 
related  to  magnesium  than  to  the  calcium  group  proper. 
Zinc  and  cadmium  are  very  much  alike,  and  are  usually 
found  together  in  nature ;  but  mercury  differs  from  them 
in  some  important  respects.     In  the  vapor  state  the  mole- 
cules of  all  three  contain  one  atom  each  {cf.  §  255). 

398.  Zinc   (Atomic    Mass,    65.4).  —  Zinc    generally 
occurs  in  combination  as  the  sulphide,  ZnS,  the  carbon- 
ate, ZnCO8,  and  the  silicate,  Zn2SiO4.     The  metallurgy 
of  zinc  is  simple.     Its  ores  are  generally  roasted  to  con- 
vert them  into  oxide,  ZnO,  and  the  oxide  is  reduced  by 
charcoal. 

The  reduction  of  zinc  oxide  takes  place  in  retorts  or  fur- 
naces; from  these  the  metal  distills  over  into  condensers.  At 
first  the  vapors  condense  as  a  powder  (zinc  dust)-,  afterwards 
the  metal  condenses  as  a  liquid,  and  is  cast  into  plates  and 
bars.  The  zinc  thus  obtained  (called  spelter)  is  not  pure. 
Pure  zinc  is  obtained  by  the  electrolysis  of  zinc  chloride. 

399.  Properties.  —  Zinc  is  a  white,  hard,  and  lustrous 
metal.     In  dry  air  it  does  not  change  ;  but  ordinarily  its 

361 


362  ZINC,    CADMIUM,    AND  MEBCURY. 

luster  is  soon  dulled  by  a  covering  of  basic  carbonate 
At  the  ordinary  temperature,  zinc  is  brittle  ;  but  betweei 
100°  C.  and  150°  C.  it  can  be  rolled  into  sheets  am 
drawn  into  wire.  At  higher  temperatures  it  is  so  brittL 
that  it  can  be  powdered. 

Zinc  melts  at  420°  C.,  and  boils  at  about  1000°  C 
When  heated  much  above  the  melting-point,  it  burn 
to  zinc  oxide.  Water  does  not  affect  the  metal,  even  a 
100°  C.  Hot  solutions  of  alkaline  hydroxides  attacl 
it,  forming  zincates  and  hydrogen  (cf.  §§  48,  336,  am 
424). 

Zn  +  2  KOH »  K2Zn02 '+  H2. 

Commercial  zinc  reacts  readily  with  all  the  ordinar 
acids ;  but  the  purer  the  metal,  the  harder  it  is  for  acid; 
to  act  upon  it. 

400.  Uses.  —  Zinc  is  used  for  the  positive  plates  o 
electric  batteries  and  in  alloys,  e.  g.,  brass  (cf.  §  409) 
Grerman  silver,  etc. 

Galvanized  iron  is  iron  coated  with  zinc  by  dipping  i 
into  a  bath  of  molten  zinc.  Galvanized  iron  resists  th< 
action  of  air  and  moisture  better  than  tinned  iron.  I 
is  used  for  wire  netting,  corrugated  roofing,  gutters 
water-tanks,  etc. 

401.  Zinc  Compounds.  —  Among  the  important  zin< 
compounds  are  the  oxide,  hydroxide,  chloride,  sulphate 
and  sulphide. 

Zinc  oxide,  ZnO,  can  be  made  by  burning  zinc,  and  by  heatin< 


PROPERTIES  AND    USES.  363 

the  carbonate  or  nitrate  (cf.  §§  169  and  321).  When-  hot  it  is 
yellow  ;  when  cold,  white.  It  is  sold  as  the  pigment,  zinc  white. 

Zinc  chloride,  ZnCl2,  may  be  formed  by  treating  the  metal 
with  hydrochloric  acid  (cf.  §  9  ).  It  is  a  white,  deliquescent 
solid  that  fuses  readily  and  distills  without  decomposing.  It 
is  usually  cast  in  sticks. 

Zinc  sulphate,  ZnSO4,  is  formed  from  the  metal  and  dilute 
sulphuric  acid.  Large  quantities  are  made  by  roasting  the 
natural  sulphide,  ZnS,  and  extracting  with  water. 

ZnS  +  2  O2  =  ZnSO4. 

From  water  the  sulphate  separates  as  transparent  crystals  of 
"  white  vitriol,"  ZnSO4.  7  II2O. 

Zinc  sulphide,  ZnS,  separates  as  a  white  precipitate  when  an 
alkaline  sulphide  is  added  to  the  solution  of  a  zinc  salt. 

402.  Cadmium.  —  The     metal    cadmium    has    the 
atomic    mass    112.3.     It    is   similar    to    zinc,   melts    at 
320°  C.,  and  is  used  in  making  some  alloys  (cf.  §  323). 
The  sulphide,  CdS,  forms  a  beautiful,  yellow  pigment. 

403.  Mercury  (Atomic   Mass,   200).  —  Mercury,    or 
quicksilver,  occurs  native  in  some  of  its  ores ;  but  the 
principal    source  of   it  is  cinnabar,  HgS.     Cinnabar  is 
mined  chiefly  in  Spain  and  California ;  recently  deposits 
have  been  found  in  Texas.     Mercury  is  obtained  from 
its  ore  by  roasting.     The  sulphur  passes  off  as  sulphur 
dioxide ;  while  the  mercury  vapors  are  condensed, 

404.  Properties   and   Uses.  —  Mercury  is   a   white, 
lustrous  liquid.      It  is  13.6  times  as  heavy  as  water; 
solidifies  at  — 39.5°  C.;  and  boils  at  357°  C. 


364  ZINC,    CADMIUM,   AND   MERCUBY. 

Mercury  does  not  react  with  hydrochloric  acid  nor 
with  cold  sulphuric  acid.  Hot,  concentrated  sulphuric 
acid  attacks  it  (cf.  §  187) ;  so  does  nitric  acid.  If 
the  nitric  acid  is  dilute  and  cold,  and  the  mercury  is  in 
excess,  the  .salt  formed  is  mercurous  nitrate,  HgNO3  ;  if 
the  concentrated  acid  is  used,  and  the  mercury  is  com- 
pletely used  up,  mercuric  nitrate,  Hg(NO3)2,  is  formed. 
Both  nitrates  are  white,  crystalline  solids. 

Mercury  vapor  is  very  poisonous. 

Mercury  is  used  in  extracting  gold  and  silver  from  their 
ores,  in  amalgamating  battery  zincs,  and  in  making  thermom- 
eters, barometers,  air-pumps,  etc. 

405.  Mercury  Compounds.  —  There  are  many  inter- 
esting mercury  compounds,  but  we  shall  consider  only 
the  oxides,  mercuric  and  inercurous,  the  chlorides,  and 
mercuric  sulphide. 

Mercuric  oxide,  HgO,  also  known  as  "  red  precipitate,"  may 
be  made  by  long  heating  of  the  metal  almost  to  its  boiling  tem- 
perature in  contact  with  air.  It  is  usually  prepared  by  heating 
the  nitrate,  Hg(NO3)2  (cf.  §  401;  zinc  oxide).  Anisomeric  form 
(cf.  §  262)  is  yellow  in  color ;  this  is  prepared  by  adding  an 
alkaline  hydroxide  solution  to  the  solution  of  a  mercuric  salt. 

2  KaOII  +  Hg(NO3)2 »  2  NaNO3  +  IlgO  +  H2O. 

The  effect  of  heat  upon  the  oxide  is  described  in  §  21. 

Mercurous  oxide,  Hg2O,  is  formed  when  an  alkali  is  added  to 
a  mercurous  salt. 

Mercurous  chloride,  HgCl,  is  known  as  calomel,  and  is  an  im- 
portant medicine.  It  is  precipitated  when  a  solution  of  a 
chloride  is  added  to  a  solution  of  a  mercurous  salt. 


MERCURY  COMPOUNDS.  365 

It  is  commonly  made  by  subliming  a  mixture  of  mercury  and 
mercuric  chloride, 


Mercuric  chloride,  HgCl2,  is  commonly  called  corrosive  sublimate. 
It  is  made  by  subliming  a  mixture  of  mercuric  sulphate  and 
sodium  chloride. 


HgSO4  +  2  NaCl  -  »  HgCl 

Mercuric  chloride  is  a  white,  crystalline  solid;  it  is  easily 
soluble  in  water,  and  very  poisonous.  It  is  used  extensively 
in  surgery  (usually  one  part  in  1,000  parts  of  water)  because  of 
its  powerful  antiseptic  action. 

Mercuric  sulphide,  HgS,  occurs  as  cinnabar  (§  403),  a  red, 
crystalline  substance.  The  sulphide  may  be  made  by  rubbing 
together  mercury  and  sulphur,  and  by  passing  hydrogen  sulphide 
into  the  solution  of  a  mercuric  salt.  In  both  cases  the  mer- 
curic sulphide  will  be  black;  but  when  it  is  sublimed  it  yields 
red  crystals.  The  red  product  is  the  pigment,  "  Chinese 
vermilion." 

Sodium  amalgam  was  mentioned  in  §§  364  and  379.  Tin 
amalgam  was  formerly  used  in  making  mirrors.  Zinc  amalgam 
is  formed  on  the  positive  plate  when  battery  zincs  are  u  amal- 
gamated." Several  amalgams  have  been  used  for  soft  fillings 
by  dentists. 


CHAPTER   XXVII. 
COPPER,    SILVER,   AND  GOLD. 

406.  Relation  of  Copper,  etc.,  to  the  Alkali  Metals, 

—  Copper,  silver,  and  gold  are  in  most  respects  differ- 
ent from  the  other  members  of  the  first  periodic  group, 
but  are  related  to  these  elements  —  the  alkali  metals  — 
much  as  zinc,  cadmium,  and  mercury  are  related  to  the 
calcium  group. 

They  are  not  changed  by  water  or  by  pure  air ;  hence 
they  occur  native  as  well  as  combined  with  other  ele- 
ments. Because  they  occur  native  all  three  have  been 
known  for  thousands  of  years,  while  none  of  the  alkali 
metals  was  known  until  1807. 

407.  Copper  (Atomic  Mass,  63.6).  — Copper  is  abun- 
dant and  widely  distributed.     It   occurs    native,    espe- 
cially in  the  Lake  Superior  region,  and  in  combination 
as    ruby   copper,    Cu2O ;    malachite,    CuCO3.  Cu(OH)2 ; 
clialcoeite,  Cu2S;  and  copper  pyrites,  Cu2S.  Fe2S3. 

Of  the  world's  supply  of  copper,  the  United  States  produces 
more  than  half  (253,870  long  tons  in  1899).  About  40%  of  this 
amount  came  from  Montana,  26%  from  Lake  Superior,  and  23% 
from  Arizona. 

Native  copper  is  obtained  by  crushing  the  ore,  washing  away 
lighter  particles,  and  smelting  and  refining  the  concentrated 

366 


USES.  367 

"mineral."  Lake  Superior  ore  has  from  0.75  of  1%  to  5%  of 
copper. 

The  Arizona  ore  consists  of  hydroxide  and  carbonate.  It  is 
smelted  with  coke  in  blast-furnaces,  and  often  yields  by  one 
fusion  a  96%  copper. 

Montana  ores  contain  sulphur  and  iron.  They  are  first  con- 
centrated by  crushing  and  washing,  and  then  roasted  to  remove 
most  of  the  sulphur.  Next  they  are  smelted  to  form  a  "  matte  " 
containing  50%  to  65%  of  copper,  besides  iron,  sulphur, 
arsenic,  etc.  ;  then  the  molten  matte  is  oxidized  in  a  "  con- 
verter" (c/.  §  432,  Fig.  63)  by  a  blast  of  hot  air,  which  re- 
moves sulphur  and  arsenic.  Finally  the  product  is  cast  into 
thick  plates  called  u  anodes,"  and  refined  by  electrolysis. 

408.  Properties.  —  Copper   has  a  red  color,  and  is 
ductile   and  malleable.     It  melts    at   about    1080°    C., 
while    silver  melts  at   954°  C.    and    gold   at    1060°  C. 
Copper  is  the  best  conductor  of  electricity  known,  ex- 
cept silver  ;  iron  is  the  only  metal  having  greater  tensile 
strength.     The  specific  gravity  of  copper  is  8.9. 

In  the  atmosphere  copper  becomes  coated  with  the 
basic  carbonate  (cf.  malachite,  §  407).  The  action  of 
nitric  acid  and  sulphuric  acid  upon  copper  has  been 
dferoSSea  in  §§  159  and  187,  respectively.  Hydrochloric 
acid  lias  practically  no  action  upon  it.  Copper  may  be 
separated  from  solutions  of  its  salts  by  zinc,  iron,  etc., 
and  by  electrolysis. 

409.  Uses.  — Copper  is  used  as  an  electric  conductor, 
as  ships'  sheathing  and  bolts,  and  for  electrical  apparatus, 
utensils,  coins,  boilers,  stills,  etc.     It  is   used,  also,  in 


368  COPPER,    SILVER,   AND   GOLD. 

copper-plating  and  electrotyping,  and  as  a  part  of  many 
alloys. 

Brass  usually  contains  28%  to  34%  zinc  and  the  remainder 
copper. 

Bronze  contains  copper,  zinc,  and  tin. 

Gun-metal  is  about  90%  copper  and  10%  tin. 

Bell-metal  consists  of  copper,  about  75%,  and  tin,  25%. 

German  silver  is,  approximately,  50%  copper,  30%  zinc,  and 
20%  nickel. 

Aluminum  bronze  is  copper  with  5%  to  10%  of  aluminum.  It 
has  the  color  of  gold,  is  hard  and  elastic,  and  does  not  tar- 
nish easily.  Aluminum  containing  3%  of  copper  is  whiter  than 
pure  aluminum. 

410.  Copper  Compounds. — Copper,  like  mercury, 
iron,  etc.,  forms  two  series  of  compounds:  cr.-prous  and 
cupric  compounds  (ef.  §§  78  and  107).  Examples  are: 

Cuprous  chloride,  CuCl  or  Cu2Cl2.      Cupric  chloride,  CuCl2. 
Cuprous  oxide,  Cu2O.  Cupric  oxide,  CuO. 

Cuprous  sulphide,  Cu2S.  Cupric  sulphide,  CuS. 

In  the  cupric  compounds  copper  is  evidently  bivalent.  If 
copper  is  bivalent  in  the  cuprous  compounds  of  the  halogens, 
e.  fj.,  in  cuprous  chloride,  the  simpler  formula  must  be  doubled. 

Cu  —  Cl. 
The  graphic  formula  for  cuprous  chloride  will  then  be    | 

Cu  — Cl. 

Cuprous  oxide  occurs  naturally  as  "ruby  copper."  It  is 
formed  as  a  black  scale  when  the  me*tal  is  heated  in  the  air, 
and  as  a  red  precipitate  when  a  solution  of  a  cupric  salt  is 
heated  with  the  solution  of  an  alkali  in  the  presence  of 
a  suitable  reducing  agent,  e.g.,  grape-sugar,  (1,;H12O6. 


COPPER-PLATING.  369 

Cupric  oxide  is  formed  by  treating  a  boiling  solution  of 
a  cupric  salt  with  the  solution  of  an  alkali.  The  blue  cupric 
hydroxide,  Cu(OH)2,  which  is  formed  if  the  solution  is  cold, 
cannot  exist  in  the  boiling  solution. 

CuS04  +  2  KOH »  K2S04  +  Cu(OH)2. 

Cu(OH)2 »  CuO  +  H20. 

Cupric  sulphate  is  known  in  crystalline  form  as  blue  vit- 
riol, the  formula  of  which  is  CuSO4.  5  H2O.  When  blue 
vitriol  is  heated,  the  water  is  expelled,  and  anhydrous 
cupric  sulphate  results.  This  is  a  white  powder  which 
becomes  blue  again  in  contact  with  water.  Blue  vitriol  is 
used  in  making  blue  and  green  pigments,  in  copper-plating, 
in  preserving  wood,  and  for  gravity  batteries. 

Cupric  sulphide,  CuS,  is  a  heavy,  black  solid,  precipitated 
when  cupric  salts  in  solution  are  treated  with  hydrogen  sul- 
phide or  with  alkaline  sulphides. 

411.  Copper-plating.  —  An  object  that  is  to  be  plated 
with  copper  is  put  into  a  bath  of  some  copper  salt  in 
solution,  and  connected  with  the  negative  electrode  (kath- 
ode) of  a  battery  or  dynamo  circuit.  A  bar  of  copper 
is  used  for  the  positive  electrode  (anode).  The  bar  of 
copper  is  gradually  used  up,  and  copper  is  deposited 
upon  the  object  to  be  plated.  The  bath  does  not  de- 
teriorate. 

This  process  is  used  in  making  electrotype  plates,  either  from 
type  or  from  woodcuts.  An  impression  of  the  type  or  wood- 
cut is  first  made  in  plaster  of  Paris  ;  this  is  covered  with 
graphite  powder  and  placed  in  the  copper-plating  bath  as  the 


370  COPPER,    SILVER,    AND   GOLD. 

kathode.     The  plate  produced  is  an  exact  reproduction  of  thn 
type  or  wood-cut  from  which  the  plaster  impression  was  taken. 

412.  Silver  (Atomic  Mass,  107.94). — Silver  is  found 
native ;  also  combined  with  sulphur  and  with  the  sul- 
phides of  other  metals.     The   lead  ore,  galena  (PbS), 
usually  contains  silver.     The  natural  chloride,  AgCl,  is 
called  "horn  silver." 

Most  of  the  world's  silver  is  found  in  the  United 
States,  Mexico,  Bolivia,  and  Australia.  The  United 
States  produced,  in  1899,  5^,764,500  ounces  —  about 
one-third  of  the  world's  supply  for  that  year. 

413.  Extraction  of  Silver  from  its  Ores.  —  Silver  is 
separated  from  its  ores  by  various  processes ;  we  shall 
consider  only  two,  viz.,  the  Amalgamation  Process  and 
the  Smelting  Process. 

The  Amalgamation  Process  consists  in  extracting  the  metal 
with  mercury,  after  a  preparatory  treatment.  This  treatment 
consists  in  crushing  the  ore,  roasting  it  with  salt  to  change  the 
sulphide  into  the  chloride  of  silver,  and  then  reducing  the 
chloride  to  silver  by  means  of  water  and  iron.  The  mass  is  then 
mixed  with  mercury,  and  the  resulting  amalgam  is  collected, 
dried,  and  heated  in  retorts.  The  mercury  distills  off,  and  is 
ready  to  be  used  again  ;  the  silver  is  then  separated  from  what- 
ever gold  is  present,  as  described  at  the  end  of  this  section. 

Smelting  Process.  —  Silver  ores  usually  contain  lead,  and  lead 
ores  often  have  enough  silver  to  pay  for  its  removal ;  hence 
the  reduction  of  the  two  metals  is  carried  out  in  a  single  smelt- 
ing operation. 

The  lead  ore  is  roasted  to  remove  sulphur,  and  then  reduced 
with  coke  in  a  blast-furnace.  The  silver  and  the  gold  present 


EXTRACTION  OF  SILVER   FROM  ITS  ORES.        371 

remain  alloyed  with  the  lead.  The  resulting  crude  lead  (known 
as  base  bullion)  is  then  heated  in  a  reverberatory  furnace  (Fig. 
60),  and  stirred  frequently.  The  small  quantities  of  copper, 


FIG.  60. 
REVERBERATORY  FURNACE. 

(The  fuel  burns  at  W;  the  substance  to  be  heated  is  placed  on  the  hearth  B. 
The  curved  roof  A  directs  the  hot  gases  of  fPdown  upon  B.) 


arsenic,  and  antimony  present  are  thus  oxidized,  and  are 
skimmed  from  the  surface  as  dross.  What  remains  is  a  mix- 
ture of  lead,  silver,  and  gold.  Most  of  the  lead  is  now  removed 
by  the  Parkes  Process.  This  consists  in  melting  the  metal  in 
large  iron  pots,  adding  1%  to  2%  of  zinc,  and  stirring.  When 
the  mixture  is  cooled  slowly,  an  alloy  of  zinc,  silver,  and  gold, 
with  but  little  lead,  comes  to  the  surface,  solidifies,  and  is 
skimmed  off.  If  necessary,  zinc  is  added  a  second  or  even 
a  third  time.  The  skimmings  are  then  heated  in  graphite  re- 
torts (c/.  §198);  and  the  zinc  is  distilled  off,  condensed,  and 
cast  into  slabs  to  be  used  again.  The  residue  in  the  retort 
consists  of  lead,  silver,  and  gold.  The  lead  is  now  completely 
removed  from  the  precious  metals  by  oxidizing  it  with  hot  air 
in  a  shallow  furnace.  The  lead  oxide  (litharge,  PbO)  flows  off 
from  the  top  of  the  furnace. 


372  COPPER,    SILVER,  AND    GOLD. 

Gold  is  separated  from  silver  by  treating  the  mixture 
of  these  metals  with  nitric  acid,  or  hot,  concentrated  sui- 
phuric  acid.  The  gold  is  not  acted  upon. 

The  sulphuric  acid  separation  is  carried  out  in  iron  kettles. 
The  silver  sulphate  formed  is  treated,  in  solution,  with  copper, 
which  precipitates  silver.  This  is  melted  and  cast  into  ingots. 

414.  Properties.  —  Silver  is  a  white  metal,  capable 
of  receiving  a  mirror-like  polish.     It  conducts  heat  and 
electricity  better  than  any  other  metal,  and  is  malleable 
and  ductile.     It  melts  at  954°  C.,  and  boils  in  the  oxy- 
hydrogen  flame.     Melted  silver  absorbs  about  twenty- 
two  times  its  own  volume  of  oxygen ;  when  the  silver 
solidifies,  the  gas  is  given  off. 

Silver  does  not  tarnish  in  pure  air,  but  is  quickly 
blackened  by  sulphur  compounds  (c/.  §  176).  Hydro- 
chloric acid  does  not  attack  it.  Nitric  acid  and  hot, 
concentrated  sulphuric  acid  act  upon  it  as  upon  copper 
(<?/.  §§  159  and  187). 

415.  Uses. — Silver  is  used    for    coinage,  tableware, 
jewelry,  ornaments,  and  mirrors,  and  for  plating  other 
metals.     Pure  silver  is  very  soft,  and  is   therefore  al- 
loyed with  copper.    The  silver  coins  of  the  United  States 
and  of  France  contain  90^>  silver  ("coin  silver"),  and 
are  said  to  be  900  fine.     The  grade  925  fine  is  called 
"  sterling  silver  "  ;  British  silver  coins  are  of  this  grade. 

Silver-plating  is  usually  done  by  electrolysis  of  the  double 
cyanide  of  silver  and  potassium.  A  bar  of  silver  forms  the 
anode,  and  the  object  to  be  plated,  the  kathode  ((/.  §  411).  The 


PHOTOGEAPHT.  373 

rough  or  "matt"  surface  is  given  the  usual  lustrous  finish  by 
polishing  with  chalk.  The  double  cyanide  solution  is  made  by 
adding  to  a  silver  nitrate  solution  one  of  potassium  cyanide  (cf. 
§  215)  until  the  silver  cyanide  first  precipitated  is  redissolved. 

Mirrors  are  made  by  precipitating  silver  upon  glass.  The 
silver  is  deposited  from  silver  nitrate  solution  containing 
ammonia.  This  solution  and  a  suitable  reducing  agent,  e.  (/., 
ammonium  tartrate,  or  acetalclehyde  (CH3CHO),  are  put  upon  the 
glass,  and  the  glass  is  gently  heated.  The  bright  deposit  of 
silver  is  washed,  dried,  and  covered  with  varnish  to  protect  it 
from  the  hydrogen  sulphide,  etc.,  of  the  air. 

416.-  Compounds  of  Silver. — Silver  nitrate,  AgNO3, 
is  the  most  important  compound  of  silver.  It  is  a  white, 
crystalline  solid,  made  by  dissolving  silver  in  dilute 
nitric  acid.  Silver  nitrate  is  sometimes  called  lunar 
caustic. 

/Silver  chloride,  AgCl,  silver  bromide,  AgBr,  and  silver 
iodide,  Agl  (cf.  §  285),  are  made  by  adding  solutions 
of  chlorides,  bromides,  and  iodides,  respectively,  to  solu- 
tions of  silver  salts.  They  are  affected  by  light,  and  are 
used  in  photography. 

417.  Photography.  —  Silver  salts  are  used  in  photo- 
graphy because  they  change  color  and  become  insoluble  in 
certain  chemicals  after  being  exposed  to  light.  When  a 
photographic  plate  is  exposed  in  a  camera,  no  change  is 
visible  until  the  plate  has  been  developed.  Developing 
consists  in  treating  the  plate  with  a  reducing  agent,  such  as 
ferrous  sulphate,  pyroyallic  acid,  hydroquinone,  eikonogen, 
etc.  When  the  plate  is  covered  with  the  developing  solu- 
tion, an  image  appears ;  this  is  due  to  the  precipitation  of  a 


374  COPPER,    SILVER,   AND   GOLD. 

film  of  silver,  which  produces  variations  of  light  and  shade. 
Where  the  light  acted  strongly  upon  the  plate,  the  deposit 
of  silver  is  relatively  heavy ;  and  where  there  was  little 
action,  there  is  little  metal  deposited. 

Just  what  the  action  of  light  upon  the  silver  bromide  of  a 
plate  is,  is  not  definitely  known,  but  it  certainly  makes  the  re- 
duction to  silver  take  place  more  easily  than  is  the  case  with 
ordinary  silver  bromide.  The  action  of  a  developer  may  be 
illustrated  by  the  following  equation  :  — 

2  AgBr  +  C6H4(OH)2  -f  2  KOH » 

hydroquinone 

2  Ag  +  2  KBr  -f-  2  H2O  -f  C6H4O2. 
quinone 

In  this  case  the  reduction  of  the  silver  bromide  is  due  to  the 
oxidation  of  hydroquinone  to  quinone. 

Fixing.  — When  the  plate  is  sufficiently  developed,  it  is  rinsed 
and  put  into  a  bath  of  sodium  thio sulphate  (u  hypo  "  ;  cf.  §  194) 
to  remove  the  silver  salts  not  acted  upon  by  light.  This  fixes 
the  negative.  Th^  plate  is  called  a  "  negative  "  because  in  it 
dark  objects  appear  light,  and  light  objects  dark. 

A  "  print "  is  made  by  placing  the  film  of  a  sensitized  paper 
next  to  the  negative  and  exposing  both  so  that  the  light  passes 
through  the  negative.  The  image  may  appear  on  the  paper  at 
once  ("printing-out"  papers)  or  may  have  to  be  deirJojx'd 
("developing"  papers).  In  either  case  the  prints  are  l*jLred" 
by  removing  the  unchanged  silver  salt. 

Toning.  —  Some  papers  are  "  toned  "  in  a  bath  of  gold,  chloride, 
AuCl3,  or  platinum  chloride,  PtCl4,  before  fixing.  Toning 
replaces  part  of  the  silver  by  gold  or  platinum. 

Blue  prints  are  made  on  paper  coated  with  a  ferric  salt  (ferric 
ammonium  citrate)  and  potassium  ferricyanide,  K8Fe(ON),.. 
After  exposure,  the  picture  is  developed  and  fixed  by  washing 


METALLURGY  OF  GOLD.  375 

it  in  waiter.     The  result  is  a  blue  print  on  a  white  ground.     The 
process  is  used  for  copying  plans,  etc. 

418.  Gold  (Atomic  Mass,  197.2). — Gold    is    found 
both  native   and    combined.     Even   native   gold  is  not 
pure,    however,    but   contains    silver,    and    often    iron, 
copper,  etc.     The  metal  is  frequently  found  enclosed  in 
quartz  or  quartz-sand. 

Gold  is  obtained  chiefly  from  Colorado  and  other  western 
states,  and  from  Australia,  Siberia,  and  South  Africa.  The 
gold  produced  in  the  United  States  during  1899  was  3,437,210 
fine  ounces,  worth  $71,053,400.  This  was  about  one-fourth  of 
the  world's  yield  that  year. 

419.  Metallurgy  of  Gold.  —  Gold-mining  is  of   two 
general  kinds :   (1)  placer-mining,  and  (2)  vein-mining. 

In  placer-mining  the  clay  and  sand  containing  the  gold  are 
washed  with  water.  The  lighter  particles  are  thus  removed; 
while  the  gold  and  other  heavy  metals  remain.  Gold  and  silver 
are  extracted  from  this  mixture  by  mercury  (r/V  §  413  ;  amal- 
gamation^. 

Vein-mining  consists  in  removing  the  gold-bearing  rock  from 
the  earth  and  crushing  it  in  stamp  mills.  The  lighter  materials 
are  then  washed  away,  and  the  gold  is  collected  with  mercury, 
as  in  placer-mining.  Hydraulic-mining  is  a  form  of  placer- 
mining  done  on  a  large  scale  with  powerful  streams  of  water. 

Instead  of  mercury,  chlorine  and  bromine  are  used  to  re- 
move gold  from  the  crushed  ore.  They  form  the  soluble  gold 
chloride,  AuCl8,  or  bromide,  AuBrs.  This  is  extracted  with 
water,  and  the  gold  is  precipitated  by  means  of  charcoal  or 
ferrous  sulphate  jftf.  §  420). 

The  Cyanide  Process  depends  upon  the  fact  that  gold   is 


376  COPPER,    SILVER,   AND    GOLD. 

converted  into  the  soluble  double  cyanide,  KCN.  AuCN,  by  a 
solution  of  the  alkali  cyanides.  The  gold  is  separated  from 
the  cyanide  solution  by  electrolysis  or  by  means  of  zinc. 

420.  Purification  of  Gold.  —  The  gold  obtained  by 
the  processes  described  above  is  not  pure.     It  can  be 
separated  from  silver  by  adding  aqua  regia,  which  reacts 
with  the  gold.      The  solution  is  evaporated  to  remove 
nitric  acid,  the  residue  is  dissolved  in  water,  and  the  gold 
is  precipitated  by  ferrous  sulphate  or  some  other  reducing 
agent. 

3  FeSO4  +  AuCl3 »  Fe2(SO4)8  -f  FeCl8  -f  Au. 

In  the  treatment  of  silver  and  gold  with  sulphuric  acid  (</.  § 
413,  end)  gold  is  left  in  the  kettle  as  a  brown,  spongy  mass. 
This  is  washed,  dried,  and  melted  in  a  crucible  with  charcoal 
and  sodium  carbonate.  The  resulting  product,  chemically  pure 
gold,  is  poured  into  a  mould,  and  leaves  it  as  a  gold  brick. 

421.  Properties  and  Uses.  —  Gold  is  the  only  common 
metal  that  is  yellow.     It  is  a  good  conductor,  and  the 
most   ductile    and    malleable    substance    known.      Its 
specific    gravity  is  19.3,  and    its    melting   temperature 
1060°  C. 

Gold  unites  directly  with  chlorine  and  bromine,  but 
not  with  oxygen.  Aqua  regia  reacts  with  gold,  forming 
auric  chloride,  AuCl3  (cf.  §  417  ;  "  toning  ") ;  but  the 
common  acids  do  not  affect  it. 

Gold  is  the  standard  of  coinage  of  most  nations.  It 
is  hardened  by  alloying  it  with  copper  (10J&  in  the 
United  States),  For  jewelry  the  proportion  of  gold 


PROPERTIES  AND    USES.  377 

varies  from  40^  to  75^ ;  it  is  usually  given  as  so  many 
"  carats  fine"  Pure  gold  is  2J,  carats  fine  ;  hence  18- 
carat  gold  is  75J&  gold  and  25^>  alloy.  Because  of  its 
malleability  and  weak  chemical  action,  gold  is  much  used 
by  dentists  for  filling  teeth.  Gold-leaf  is  used  in  or- 
namentation. 


CHAPTER   XXVIII. 
ALUMINUM  (Atomic  Mass,  27). 

422.  Occurrence  of  Aluminum. — Although  alumi- 
num does  not  occur  free,  it  is   the   most  abundant  and 
widely  distributed  metal.     Only  oxygen  and  silicon  are 
more  abundant.     Some  of  the  most  important  minerals 
containing    aluminum    are  feldspar    (KAlSi3O8),    mica 
(KAlSiO4),  and  cryolite  (cf.  §  266).      Granite  is  a  mix- 
ture of  quartz,  feldspar,  and  mica.      Clay  results  when 
granite  and  similar  minerals  are  decomposed. 

All  the  other  elements  of  the  aluminum  group  are 
rare  (c/.  Periodic  Table,  §  332). 

423.  Preparation.  —  Aluminum   was    formerly   pro- 
duced from  anhydrous  aluminum  chloride  and  sodium.* 

A1C18  +  3  Na »  Al  -f  3  NaCl. 

This  method  has  been  superseded  by  electrolytic 
processes,  of  which  that  of  Hall  (1887)  is  perhaps  the 
most  important. 

Hall's  Process. — The  furnace  used  in  the  Hall  process  is  a 
box  of  boiler  iron  (Fig.  61),  the  bottom  and  sides  of  which  are 
lined  with  a  mixture  of  coke  and  tar,  rammed  hard.  The  bot- 
tom forms  the  —  electrode,  while  the  +  electrode  consists  of 
forty  large  carbons  suspended  by  copper  rods. 

To  begin  the  process,  the  carbons  are  lowered  almost  to  the 

878 


PROPERTIES. 


379 


bottom  of  the  furnace,  cryolite  is  put  in,  and  the  current  is 
turned  on.  The  resistance  to  the  passage  of  the  current  pro- 
duces heat  enough  to  melt  the  cryolite.  Pure,  dry  aluminum 


FIG.  61. 

oxide,  A12O3,  is  then  mixed  with  the  fused  cryolite,  and  the 
electrolysis  begins. 

The  process  is  continuous,  for  the  cryolite  bath  remains  un- 
changed. The  aluminum  collects  at  the  bottom,  and  is  drawn 
off  ;  the  oxygen  of  the  aluminum  oxide  unites  with  the  carbons 
to  form  carbon  monoxide,  which  escapes.  One  company  getting 
its  power  from  Niagara  produced  3,190  tons  of  aluminum  dur- 
ing 1900  ;  this  was  about  one-third  of  the  world's  output. 

The  aluminum  oxide  used  is  obtained  from  beauxite  (or 
bauxite),  A1(OH)3.  The  natural  mineral  has  impurities,  e.  g., 
iron,  silicon,  etc.  ;  these  are  removed  at  present  by  fusing  the 
beauxite  with  a  little  metallic  aluminum. 

424.  Properties.  —  Aluminum  is  a  white,  lustrous 
metal.  Its  specific  gravity  (2.6)  is  very  low  compared 
with  that  of  other  common  metals,  zinc  being  7.1  and 


380  ALUMINUM. 

iron  7.8  times  as  heavy  as  water.  It  melts  at  about 
660°  C.  and  vaporizes  at  the  temperature  of  the  electric 
furnace.  Aluminum  is  a  good  conductor,  is  ductile 
and  malleable,  and  has  the  tensile  strength  of  cast- 
iron.  Commercial  aluminum  is  95^  to  99.  6/0  pure. 

At  white  heat  aluminum  burns  to  aluminum  oxide, 
A12O3.  Hydrochloric  acid  readily  reacts  with  the  metal, 
forming  aluminum  chloride. 


2  Al  +  6  HC1  -  »  2  AlCl3  +  3  Ha. 

Nitric  acid  and  dilute  sulphuric  acid  do  not  act  upon  it 
ordinarily. 

Aluminum  reacts  with  solutions  of  salt  and  other 
chlorides  if  a  little  free  acid  is  present.  It  reacts, 
also,  with  the  hydroxides  of  sodium  and  potassium,  form- 
ing aluminates  and  hydrogen. 

6  NaOH  +  2  Al  -  >  2  ]STa3AlO3  -f  3  H2. 

Aluminum  unites  directly  with  the  halogens,  and  with 
carbon,  silicon,  nitrogen,  etc. 

425.  Uses.  —  It  is  probable  that  more  aluminum  is 
used  in  the  manufacture  of  iron  and  steel  than  for  any 
other  one  purpose.  The  aluminum  removes  any  oxygen 
the  iron  may  have  in  combination,  and  thus  increases 
the  fluidity  of  cast-iron  and  steel. 

The  next  important  use  is  as  a  conductor  of  elec- 
tricity. The  Northwestern  Elevated  Railroad  of  Chi- 
cago is  using  twenty  miles  of  1^-inch  aluminum  cables, 


ALUMINUM  OXIDE  AND  HYDROXIDE.  381 

weighing  150,000  pounds,  to  transmit  motive  power  to 
its  trolley  cars. 

Aluminum  powder  is  used  to  reduce  the  oxides  of  many 
metals,  e.  g.,  chromic  oxide,  Cr2O3,  and  for  flash-lights.  A  large 
part  of  the  aluminum  produced  is  used  for  kitchen  utensils, 
scientific  instruments,  etc.  ;  the  aluminum  alloys  also  require 
much  of  the  metal. 

The  alloys  with  copper  were  described  in  §  409.  A  new 
alloy,  called  magnalium,  contains  75%  to  90%  aluminum  and  the 
remainder  magnesium. 

426.  Aluminum  Oxide  and  Hydroxide.  —  Alumi- 
num oxide,  A12O3,  occurs  in  the  form  of  ruby,  sapphire, 
and  corundum.  Emery,  an  impure  form  of  corundum, 
is  very  hard,  and  is  used  for  grinding  and  polishing. 
Aluminum  oxide  may  be  made  by  heating  the  hydroxide. 

2  A1(OH)8  =  A12O3  +  3  H2O. 

Aluminum  hydroxide,  A1(OH)3,  may  be  precipitated 
by  adding  ammonium  hydroxide  to  the  solution  of  an 
aluminum  salt,  e.  g.,  aluminum  chloride.  Aluminum 
hydroxide  reacts  with  both  acids  and  alkalies  (except 
ammonium  hydroxide)  ;  with  acids  it  gives  aluminum 
salts,  and  with  alkalies,  aluminates  (cf.  §  424).  It  is, 
therefore,  a  base  toward  strong  acids  and  an  acid  toward 
strong  bases  (cf.  §§  312  and  317). 

(1)  A1(OH)8  -f      3  HC1 >  A1C18       -f  3  H2O. 

(2)  H3A1O3    -f  '*  NaOH »  Ka3AlO3  +  3  H2O. 

Aluminum  hydroxide  unites  either  chemically  or  mechani- 
cally with  many  dye-stuffs,  and  also  with  certain  fabrics.  Ad- 


ALUMINUM. 


vantage  is  taken  of  this  fact  to  "  fix  "  the  dyes  in  fabrics  that 
do  not  readily  hold  color.  The  hydroxide  is  therefore  called 
a  mordant,  from  Latin  to  bite,  because  it  bites  into  the  fabric. 

427.  Aluminum  Salts.  —  Aluminum  chloride,  A1C13, 
is  soluble  in  water,  and  crystallizes  with  crystal- 
water  (AlClg.  6  H2O)  ;  but  it  is  so  much  hydrolyzed 
that  when  the  water  is  expelled  hydrochloric  acid  es- 
capes, and  aluminum  hydroxide  remains. 

A1C18  +  3  H20  -  -  »  A1(OH)3  +  3  HC1. 

The  anhydrous  chloride  is  made  by  letting  dry  hydro- 
chloric acid  gas  act  upon  hot  aluminum  filings.  It  is  a 
hygroscopic,  white  powder. 

Alums.  —  Aluminum  sulphate,  A12(SO4)3,  forms  double  salts 
with  the  sulphates  of  univalent  metals,  e.  g.,  with  potassium 
sulphate,  K2SO4  ;  these  double  sulphates  are  called  alums.  Tho 
alums  crystallize  with  twenty-four  molecules  of  crystal-water. 
Potash-alum  is  K2SO4,  A12(SO4)3,  24  H2O  ;  silver-alum  is 
Ag2SO4,  A12(SO4)3,  24  H2O.  Other  trivalent  elements  may 
replace  aluminum.  Thus  N"a2SO4,  Ee2(SO4)3,  24  H2O  would  be 
sodium  iron-alum.  Chrome-alum  is  K2SO4,  Cr2(SO4)3,  24  H2O. 

Aluminum  Carbonate  and  Sulphide.  —  The  elec- 
tro-positive properties  of  aluminum  are  so  weak  that 
even  its  salts  with  strong  acids,  e.  g.,  the  chloride  and 
sulphate,  are  readily  hydrolyzed.  Its  salts  with  weak 
acids  cannot  exist  in  the  presence  of  water,  being  com- 
pletely decomposed  into  the  hydroxide  and  the  free  acid. 

When  the  acid  is  volatile,  it  of  course  escapes.  This  is  the 
case  with  the  carbonate.  When  aluminum  «alts  are  treated 


PORCELAIN,    STONEWARE,   ETC.  383 

with  the  solution  of  a  carbonate,  the  products  are  aluminum 
hydroxide  and  carbonic  acid.     Hence  carbon  dioxide  escapes. 

(1)  2  A1C13      +  3  N"aoCO3 »  A10(CO3)3  -f  6  NaCl. 

(2)  A12(C03)3  +  G  H26       »  2  A1(OII)3  -f  3  H2CO3. 

With  the  sulphide  the  result  is  similar. 

(1)  2  A1C13  +  3  (NII4)2S »  A12S3          +  6  NH4C1. 

(2)  A12S3     -f  6  H2O         »  2  A1(OII)8  -f  3  H2S. 

428.  Porcelain,  Stoneware,  Etc.  —  Aluminum  sili- 
cate, Al2(SiO8)3,  is  essentially  the  substance  from  which 
porcelain,  etc.,  are  made.  In  a  pure  form  it  is  kaolin /  in 
an  impure  form,  day  (cf.  §  341). 

Porcelain  is  made  by  mixing  white  kaolin  with  more  fu- 
sible substances,  such  as  feldspar,  shaping  the  plastic  mixture 
into  form,  and  heating  it  to  a  high  temperature.  The  more 
fusible  portion  (the  feldspar)  melts,  and  cements  the  whole 
together.  Porcelain  is  hard  and  translucent,  and  withstands 
the  action  of  heat  and  chemicals  better  than  glass,  hence  it 
is  used  for  many  purposes  in  chemical  laboratories. 

Stoneware  is  opaque,  for  it  has  not  been  heated  enough 
to  make  the  feldspar  penetrate  the  kaolin  as  much  as  in 
porcelain. 

}£arthenware  is  made  from  common  clay,  hardened  by 
heat,  but  not  fused.  It  is  glazed  by  putting  common  salt 
into  the  furnace  at  the  time  of  heating.  This  forms  a 
covering  of  sodium  aluminum  silicate  over  the  porous  sur- 
face. Bricks,  tiling,  jugs,  terra- cotta,  etc.,  are  examples. 

Ultramarine  is  a  blue  coloring  substance,  made  by  melting 
together  kaolin,  sodium  carbonate,  and  sulphur.  This  substance 
was  once  very  valuable,  but  thousands  of  tons  of  it  are  now 
made  every  year. 


CHAPTER   XXIX. 
IRON,  COBALT,  AND  NICKEL. 

429.  Iron  (Atomic  Mass,  56).  — Iron  is  the  most  use- 
ful of  the  metals.     It  is  also  one  of  the  most  widely 
distributed,  since  it  is  found  in  many  minerals,  in  the 
soil,  and  in  natural  waters.     Iron  is  an  essential  part  of 
chlorophyll,  the  green  material  of  plants,  and  of  the  red 
coloring  matter  of  the  blood.     It  is  present  in  meteor- 
ites, and  in  the  sun  and  stars. 

The  principal  ores  of  iron  are  hcematite  (Fe2O3), 
magnetite  (Fe3O4),  broivn  iron  ore  [Fe2O3.  2  Fe(OH)3], 
and  siderite,  or  spathic  iron,  (FeCO3).  Iron  pyrites, 
FeS2,  is  a  source  of  sulphur. 

430.  Metallurgy.  —  The  ores  of  iron  are  reduced  by 
heating  them  with  carbon  (coke  or  coal)  in  a  blast-fur- 
nace   (Fig.    62).     A  flux   (limestone    or    feldspar)    is 
added  to  combine  with  the  ashes  of  the  coal  and  form 
a  slag  (<?/.  §  266). 

A  blast-furnace  (Fig.  62)  is  a  structure  from  thirty  to  ninety 
feet  high.  The  inner  walls  are  of  fire-brick  surrounded  by 
brick  or  stone  ;  the  outside  is  made  of  sheet-iron.  The  fur- 
nace is  nearly  filled  from  the  top  (A)  with  successive  layers  of 
fuel,  ore,  and  flux  ;  while  a  blast  of  hot  air  is  forced  in  through 
pipes  (tuyeres)  at  the  bottom.  The  reduced  iron  collects  as  a 
liquid  at  the  bottom  of  the  furnace,  and  above  it  the  slag. 

384 


METALLURGY. 


385 


More  fuel,  ore,  and  flux  are  added  frequently,  so  that  the  fur- 
nace is  kept  filled.     The  operations  require  careful  attention. 
The  ores  must  be  analyzed,  the 
nature   and   amount  of   the  flux 
determined,  and  the  temperature 
controlled.     The    iron    and    the 
slag  are  drawn  off   through   the    C 
tap-holes   I  and  S,  respectively, 
two  or  three  times  a  day.     The 
iron   is   run   into  moulds,  form- 
ing   the    bars    called    "pigs" 
(hence  the  name  pig-iron),  or  it 
is  transferred   to   the    converters 
((/.  §  432)  and  made  into  steel. 
Cast-iron  is  a  form  of  pig-iron. 

When  once  started,  blast-fur- 
naces continue  in  operation  for 
months.  The  largest  ones  yield 
five  hundred,  or  more,  tons  of 
pig-iron  a  day. 


FIG. 


The  chemical  reactions  in- 
volved are  as  follows :  The 
oxygen  of  the  air-blast  unites 

with  the  carbon  of  the  fuel,  in  the  lower  part  of  the 
furnace,  to  form  carbon  dioxide  ;  a  little  higher  up,  this 
is  reduced  by  the  hot  fuel  to  carbon  monoxide.  It  is 
carbon  monoxide  that  reduces  the  ore. 


Fe0 


2  Fe  +  3  C02. 


The  waste  gas-es  escaping  from  the  furnace  at  C  are  about 
25%  carbon  monoxide;  this  is  burned  to  produce  steam  for  the 
engines  operating  the  blast  or  doing  other  work. 


386  IE  ox,  con  ALT,  AND  NICKEL. 

431.  Commercial  Iron.  —  The  various  grades  of  iron 
are  alloys  of  the  metal  with  more  or  less  carbon,  besides 
traces  of  silicon,  sulphur,  phosphorus,  manganese,  etc. 
The  three  chief  subdivisions  of  the  kinds  of  iron  are 
cast-iron,  wr  ought-iron,  and  steel. 

Cast-iron  contains  from  1.5%  to  7%  of  carbon.  All  the 
varieties  of  cast-iron  melt  comparatively  low  (1050°  to 
1250°  C.),  and  are  shaped  by  pouring  them  in  the  liquid 
state  into  sand  moulds.  Cast-iron  is  too  brittle  to  be  welded 
or  forged. 

Wr ought-iron.  —  If  the  carbon  of  cast-iron  is  nearly  all  re- 
moved, the  iron  becomes  tough  and  malleable,  and  requires  a 
high  temperature  (1400°  C.  and  above)  to  melt  it.  It  can  be 
forged  and  welded,  but  not  cast  or  tempered.  This  form  is 
wr ought-iron.  It  usually  contains  0.6%,  or  less,  of  carbon. 

Steel  generally  contains  more  carbon  than  wrought- iron, 
and  less  than  cast-iron.  It  may  be  forged,  welded,  cast, 
and  tempered.  It  melts  higher  than  cast-iron  and  lower 
than  wrought-iron. 

Annealing  and  Tempering  Steel.  —  If  steel  containing  over 
0.5%  of  carbon  is  heated  to  cherry  redness  and  then  quickly 
cooled  in  water  or  oil,  it  hardens,  and  is  suitable  for  cutting 
tools,  etc.  If  the  hardened  steel  is  slowly  heated  up  again, 
and  then  slowly  cooled,  it  becomes  soft.  This  process  is  known 
as  annealing.  If,  however,  the  hardened  steel  is  heated  up 
slowly  until  a  film  of  a  particular  color  appears,  and  is  then 
plunged  into  water,  it  will  retain  a  definite  degree  of  hardness. 
This  process  is  called  tempering. 

432.  Manufacture  of  Steel. — Steel  may  be  made 
(1)  by  removing  part  of  the  carbon  from  cast-iron,  (2) 


MANUFACTURE   OF  STEEL. 


387 


by  adding  carbon  to  wrought-iron,  and  (3)  by  melting 
together  cast-  and  wrought-iron.  Method  (1)  is  now 
rarely  used;  method  (2)  is  applied  in  the  crucible  pro- 
cess of  making  steel ;  method  (3)  is  the  basis  of  the 
Bessemer  and  Open  Hearth  (Siemens-Martin)  processes. 

The  crucible  process  consists  essentially  in  heating  a  very 
pure  wrought-iron  with  carbon  for  a  long  time.  Some  of  the 
carbon  is  absorbed,  producing  a  very  fine  quality  of  steel,  suit- 
able for  tools.  The  process  is,  however,  expensive. 

The  Bessemer  process  consists  essentially  in  reducing  pig 
iron  in  a  "converter"  to  wrought-iron,  and  then  adding 
enough  cast-iron  (called  "  spiegeleisen")  to  bring  the  propor- 
tion of  carbon  up  to  the  desired  point. 

The  converter  (Fig.  63)  is  a  large,  pear-shaped  furnace 
mounted  upon  supports  (8)  so  that  it  can  be  inverted.  A 
blast  of  air  forced  up  through  openings 
at  the  bottom  (tuyeres,  T)  oxidizes  the 
silicon  and  carbon,  the  former  to  a  slag, 
the  latter  to  carbon  dioxide.  Manganese 
is  put  into  the  spiegeleisen  to  reduce  any 
iron  oxide  formed,  the  resulting  oxide  of 
manganese  separating  in  the  slag.  If  the 
pig-iron  contains  much  phosphorus  and 
sulphur  the  converter  is  lined  with  quick- 
lime and  magnesia.  The  lining  retains 
most  of  the  phosphorus  as  phosphate. 

Converters  usually  make  ten  to  twenty 
tons   of    steel   at  a   "  blow."      The   time  FIG. 

taken  is  about  thirty  minutes. 

Bessemer  steel  contains  about  0.35%  carbon  ;  it  is  used  for 
rails,  axles,  cannon,  wire,  tin-plate,  and  structural  purposes. 

In  the  Open  Hearth  process,  pig-iron  is  mixed  with  wrought- 


388  IRON,    COBALT,   AND  NICKEL. 

iron  or  steel  scrap,  and  heated  on  a  hearth  with  an  oxidizing 
gas-flame.  When  the  proportion  of  carbon  has  been  lowered 
sufficiently,  manganese  is  added  in  the  form  of  "  spiegeleisen" 
or  "  ferro-manganese."  The  charge  is  usually  about  twenty 
tons;  but  the  process  takes  from  eight  to  eleven  hours,  and  is 
consequently  more  expensive  than  the  Bessemer  process. 
Steel  made  in  this  way  is,  however,  very  tough  and  elastic,  and 
is  suitable  for  the  finest  structural  work,  e.  g.,  bridges,  and  for 
machinery,  boiler-plate,  large  guns,  etc. 

433.  Properties   of   Iron.  —  Pure   iron   is   rare ;   it 
melts  at  about  1800°  C.     The  purest  commercial  form 
is  wrought-iron ;   this  is  malleable,  ductile,  and  a  fairly 
good    conductor.     In    the   form  of   steel,  iron   may  be 
made  very  hard. 

Soft,  i.  e.,  wrought,  iron  may  be  magnetized  temporarily,  but 
soon  loses  the  magnetism  ;  steel  is  not  so  easily  magnetized, 
but  becomes  a  permanent  magnet. 

At  a  high  temperature  iron  decomposes  water  vapor, 
yielding  the  oxide  Fe3O4  and  hydrogen. 

3  Fe  -f  4  H2O »  Fe3O4  -f  4  H2. 

When  iron  is  burned  in  oxygen  the  same  oxide  re- 
sults (cf.  §  23).  In  dry  air  iron  does  not  change,  but 
in  the  presence  of  moisture  and  carbon  dioxide  it  rusts. 
Iron  rust  is  a  mixture  of  ferric  oxide  (Fe2O3)  and  ferric 
hydroxide,  Fe(OH)3.  Iron  reacts  easily  with  dilute 
acids.  Two  classes  of  iron  compounds  are  known,  viz., 
ferrous  and  ferric  compounds  (cf.  J§  78  and  107). 

434.  Oxides    and    Hydroxides  of    Iron.  —  Ferrous 
oxide,  FeO,  is  not  easily  prepared  or  kept  in  pure  condition. 


IRON  CHLORI'DES.  389 

It  may  be  made  by  reducing  ferric  oxide  with  hydrogen  at 
300°  C.  ;  but  on  exposure  to  air  it  at  once  oxidizes. 

ferric  oxide,  Fe2O8,  is  found  in  enormous  masses 
(haematite).  It  may  be  made  by  heating  ferric  hydrox- 
ide or  ferrous  sulphate. 


Ferrous-ferric  oxide,  Fe3O4,  is  called  the  "magnetic  ox- 
ide "  of  iron  ;  it  occurs  as  magnetite.  It  is  sometimes  found 
as  lodestone,  a  natural  magnet.  Iron  may  be  kept  from 
rusting  by  exposing  it  while  red  hot  to  steam  ;  the  resulting 
layer  of  the  black  oxide,  Fe3O4,  protects  the  remainder  of 
the  metal. 

ferrous  hydroxide,  Fe(OH)2,  is  formed  when  an  alkali 
in  solution  is  added  to  the  solution  of  a  ferrous  salt.  It  is 
usually  green,  but  soon  becomes  brown  where  air  is  in  con- 
tact with  it.  In  the  absence  of  air  it  is  white. 

Ferric  hydroxide,  Fe(OH)3,  is  made  by  adding  an  al- 
kali or  ammonia  water  to  the  solution  of  a  ferric  salt.  It 
forms  a  flaky,  red-brown  precipitate. 

435.  Iron     Sulphides.  --  Ferrous     sulphide,     FeS     (c/. 
§  179),  is  a  black  solid  made  by  heating  a  mixture  of  sulphur 
and  iron,  and  by  adding  an  alkali  sulphide   to   a  ferrous   salt 
solution. 

K2S  -f  FeCl2  -  »  FeS  +  2  KC1. 

Iron  pyrites  (FeS2)  has  the  color  of  brass,  and  is  called 
"fool's  gold."  For  its  behavior  when  heated,  see  §  178. 

436.  Iron    Chlorides.  —  Ferrous   chloride,    FeCl2,    is 
made  by  passing  hydrochloric  acid  gas  over  hot  iron. 
A  solution  of  it  results  when  iron  is  treated  with  the 


390  IRON,    COBALT,    AND  NICKEL. 

acid.     Like  all  ferrous  compounds  it  oxidizes  easily  to 
the  ferric  condition. 

In  the  presence  of  hydrochloric  acid  the  oxidation  produces 
ferric  chloride. 

4  FeCl2  +  4  IIC1  +  02 >  4  FeCl8  +  2  H2O. 

If  no  acid  is  present,  part  of  the  iron  is  oxidized  to  ferric 
salt,  and  part  to  ferric  hydroxide  (rust). 

12  FeCl2  +  3  02  +  6  H2O >  8  FeCl3  +  4  Fe(OH)8. 

Ferric  chloride,  FeCl3,  is  formed  in  solution  by  passing 
chlorine  into  ferrous  chloride  solution,  or  by  treating  iron 
with  aqua  regia,  and  evaporating  repeatedly  with  hydro- 
chloric acid.  The  anhydrous  salt  is  made  by  passing 
chlorine  over  red-hot  iron.  It  looks  like  maple-sugar. 

437.  Iron   Sulphates.  —  A    solution    of  ferrous  sul- 
phate, FeSO4,  results  when  iron  reacts  with  dilute  sul- 
phuric  acid.     The    crystals  known    as   green  vitriol  or 
"  copperas  "  are  FeSO4.  7  H2O.     Green  vitriol  is  used 
in  making  inks,  in  dyeing,  and  as  a  deodorizer.      Fer- 
rous   sulphate    is    oxidized    like    the    chloride.      With 
ammonium    sulphate  it   forms    the   double  salt,  ferrous 
ammonium  sulphate,  (NH4)2SO4.  FeSO4.  6  H9O ;    this  is 
much  less  easily  oxidized  than  ferrous  sulphate  alone. 

Ferric  sulphate,  Fe2(SO4)3,  is  made,  in  solution,  by 
oxidizing  ferrous  sulphate  with  nitric  acid  in  the  pres- 
ence of  sulphuric  acid. 

438.  Potassium  Ferro-  and  Ferri-cyanides.  —  Potas- 
shim  ferrocyanide,  K4Fe(CN)6,  is  called  "yellow  prussiate" 


COBALT.  391 

(i.  e.,  cyanide,  cf.  §  216)  "of  potassium."  It  is  a  yellow, 
crystalline  solid.  With  a  ferric  salt  it  produces  Prussian 
blue.  Potassium  ferricyanide,  K3Fe(CN)6,  is  obtained  by 
oxidizing  the  ferrocyanide  with  chlorine.  It  is  a  red,  crys- 
talline solid. 

Potassium  /e?vocyaiiide  may  be  written  Fe(CN)2. 
4  KON  ;  this  formula  shows  that  the  iron  is  ferrous,  i.  e., 
bivalent.  (Ferro-  means  ferrous.)  In  the  /err/cyanide, 
Fe(CN)3.  3  KCN,  iron  is  in  the/erHc  condition. 

439.  Nickel    (Atomic  Mass,  58.7).  —  Nickel    occurs 
with  iron  in  meteorites.     Its  ores  (silicates  of  nickel)  are 
found  chiefly  in  Canada,  Norway,  and  New  Caledonia. 
Like  iron,  it  forms  two  classes  of  compounds ;  these  are 
nickelous  and  nickelic  compounds.     The  common  com- 
pounds are  nickelous.     Most  of  them  are  green.     The 
formula    of   nickel  sulphate  is   NiSO4 ;    that   of  the  ni- 
trate, Ni(NO3)2;  that  of  the  sulphide,  NiS. 

Nickel  is  used  to  plate  other  metals  to  protect  them 
from  the  atmosphere.  It  is  used,  also,  in  making  alloys, 
e.  g.,  German  silver,  nickel  steel,  and  the  United  States 
five-cent  piece.  Both  nickel  and  cobalt  are  attracted  by 
the  magnet. 

440.  Cobalt   (Atomic  Mass,  59).  —  Cobalt  occurs  com- 
bined  with    arsenic   and   sulphur,  and  often    associated   with 
nickel.     Cobalt  salts  are  red  in  solution  or  combined  with  much 
crystal-water,  and  blue  when  anhydrous  or  with  little  crystal- 
water.     A  solution  of  cobaltous  chloride,  CoCl2,  is  used  as  a 
sympathetic  ink. 


CHAPTER   XXX. 
MANGANESE   AND   CHROMIUM. 

441.  Manganese   (Atomic   Mass,   55).  —  Manganese 
occurs  chiefly  as  the  black  oxide,  MnO2,  —  the  mineral 
pyrolusite.     The  pure  metal  is  very  hard,  and  fuses  only 
at  a  high  temperature.     In  some  ways  it  resembles  iron. 
Thus,  it  forms  two  series  of  salts,  manganous  and  man- 
ganic  salts,   corresponding  to  ferrous   and  ferric  salts. 
The  manganous  salts,  however,  are  more  stable  than  the 
ferrous    salts,  and  are  not  readily  oxidized.     Manganic 
salts  are  much  less  stable  than  manganous  salts.     The 
latter  are  usually  pink  in  color  and  crystalline.     They 
are  formed  when  the  higher  oxides  of  manganese  are 
treated  with  acids,  oxygen  being  either  set  free  or  else 
used  in  oxidizing  the  acid  (ef.  §  81). 

442.  Manganese  Oxides.  —  Manganese  forms  the  fol- 
lowing oxides :  — 

Manganous  oxide,  MnO  ;  Manganese  dioxide,  MnO2 ; 

Manganic  oxide,  Mn2O3 ;  Manganese  heptoxide,  Mn2O7. 

Manganous-manganic  oxide,  Mn3O4 ; 

All  of  these  are  solids  except  the  last,  which  is  a  dark 
liquid.  The  oxide  Mn2O5,  corresponding  to  the  mangan- 
ates,  is  not  known. 

The  most   important  oxide  of  manganese  is  the  di- 


POTASSIUM  PERMANGANATE.  393 

oxide.  This  is  used  in  preparing  chlorine  and  oxygen, 
and  in  decolorizing  glass. 

The  manganese  dioxide  used  in  making  chlorine  is  recovered 
by  treating  the  manganous  chloride  produced  (cf.  §  81)  with 
slaked  lime.  This  forms  manganous  hydroxide,  Mn(OH)2. 
By  means  of  steam,  air,  and  more  lime,  this  is  converted  into 
manganites  having  the  formulas  CaMnO3  (i.  e.,  CaO.  MnO2) 
and  CaMn2O5  (i.  e.,  CaO.  2  MnO2).  Both  of  these  give  chlorine 
when  treated  with  hydrochloric  acid. 

443.  Potassium  Permanganate.  —  Manganese  forms 
not  only  salts,  in  which  the  manganese  is  the  electro- 
positive element,  but  also  manganates  and  permangan- 
ates, in  which  the  manganese  has  the  same  relation  to 
the  compound  that  sulphur  has  to  the  sulphates.  The 
more  highly  oxidized  the  manganese  is,  the  less  basic 
does  its  oxide  become.  Manganese  heptoxide,  Mn2O7,  is 
the  anhydride  of  permanganic  acid,  HMnO4;  the  most 
important  salt  of  this  acid  is  potassium  permanganate, 
KMnO4. 

Potassium  permanganate  is  formed  by  boiling  the 
solution  of  the  manganate,  K2MnO4,  and  passing  carbon 
dioxide  or  chlorine  into  it. 

(1)  3  K2Mn04  +  2  CO2  =  2  K2CO3  +  2  KMnO4  +  MnO2. 

(2)  2  K2MnO4  +  C12  =  2  KC1  +  2  KMnO4. 

The  permanganate  separates  from  solution  in  prisms.  It  col- 
ors water  a  deep  purple.  The  manganate  is  obtained  by  fusing 
a  mixture  of  manganese  dioxide,  potassium  hydroxide,  and  an 
oxidizing  agent,  e.  g.,  potassium  nitrate  or  chlorate. 

A  crude  permanganate  solution  is  used  to  oxidize  sewage. 


394  MANGANESE  AND    CHROMIUM. 

it  is  made  by  treating  crude  sodium  manganate  or  potassium 
manganate  with  dilute  sulphuric  acid. 

444.  Oxidation  by  Permanganate.  —  Most  of  the 
uses  of  potassium  permanganate  are  due  to  its  easy 
liberation  of  oxygen.  There  is  a  difference  depending 
upon  whether  it  acts  in  acid  or  in  alkaline  solution. 

We  represent  the  oxidation  of  oxalic  acid  by  potas- 
sium permanganate  in  the  presence  of  sulphuric  acid  as" 
follows  :  - 

2  KMnO4  -f  5  H2C2O4  +  3  H2SO4  =  2  Mn$O4  +  K2$O4  -f 


Another  process  depends  upon  the  oxidation  of  a 
ferrous  to  a  ferric  compound. 

2  KMn04  +  10  FeS04  +  8  H2SO4  ==  5  Fe2(SO4)3  +  K2SO4  + 
2  MnSO4  +  8  H2O. 

Botli  of  these  reactions  are  sharp  and  complete,  hence 
they  have  important  uses  in  volumetric  analysis. 

To  understand  the  oxidizing  action  of  potassium  perman- 
ganate in  acid  solution  we  must  look  upon  this  substance  as 
made  up  of  potassium  oxide  and  manganese  heptoxide. 

2  KMnO4  =  K2O  +  Mn,O7. 

The  molecule  of  manganese  heptoxide  gives  up  five  atoms  of 

oxygen. 

Mn2O7  =  2  MnO  -f  5  O. 

The  manganmts  oxide  then  reacts  with  the  acid  to  form  a  man- 
ganous  salt  and  water,  and  the  potassium  oxide  to  form  a  po- 
tassium salt  and  water. 


CHROMIUM.  395 

MnO  +  H2S04  =  MnS04  +  H2O. 
K2O  +  H2SO4  =  K2SO4  -f  HoO. 

The  complete  equation  is,  therefore, — • 

2  KMnO4  -f  3  H2SO4  =  2  MnSO4  -\-  K2SO4  +  3  H2O  -f  5  O. 

When  the  permanganate  is  used  in  neutral  or  alkaline 
solution,  the  available  oxygen  is  less  than  in  acid  solu- 
tion, e.  y.,  - 

2  KMnO4  -f  H2O  =  3  O  +  2  KOH  -f  2  MnO2. 

If  there  is  sufficient  alkali,  the  manganese  dioxide  unites 
with  it  to  form  a  manganite,  e.  g.,  K2MnO3,  which  re- 
mains in  solution. 

The  action  of  potassium  permanganate  with  sulphurous  acid 
and  with  hydrogen  peroxide  is  represented  thus  :  — 

2  KMn04  -f  5  II28O3  +  3  H2SO4  =  K2SO4  +  2  MnSO4  + 

5H2SO4-f3H2O. 

2  KMnO4  +  5  H2O2  +  3  H2SO4  ="K2SO4  +  2  MnSO4  + 
8  1I2O  +  5  C)2. 

445.  Chromium  (Atomic  Mass,  52.1).  —  Chromium 
is  a  comparatively  rare  element.  It  occurs  chiefly  as 
chromite,  Fe(CrO2)2  or  FeO.  Cr2O3.  The  name  of  the 
element  is  from  the  Greek  chroma,  meaning  color. 
Chromium  compounds  have  many  decided  colors. 

The  metal  is  prepared  at  a  high  temperature  by  re- 
ducing the  oxide  with  carbon  or  the  chloride  with 
sodium.  It  is  steel-gray,  very  hard,  and  very  difficult 
to  fuse.  With  iron  it  forms  a  hard  alloy  called  "chrome- 
steel." 


396  MANGANESE  AND   CHROMIUM. 

446.  Oxides  and  Hydroxides.  —  The  most  important 
oxides  of    chromium    are    chromic    oxide    (Cr2O3)    and 
chromium  trioxide  (CrO3).      Chromous  oxide,  CrO,  is  the 
one  from  which  the  chromous  salts  are  derived. 

Chromic  oxide  is  a  valuable  pigment  known  as  "  chrome- 
green."  It  is  made  by  driving  off  water  from  the  hydroxide, 
Cr(OH)3,  or  by  heating  a  mixture  of  potassium  dichromate, 
ammonium  qhloride,  and  sodium  carbonate. 

Chromium  trioxide  separates  as  bright  red  crystals  when 
strong  sulphuric  acid  is  added  to  a  saturated  solution  of  potas- 
sium dichromate. 

Although  often  called  "  chromic  acid,"  it  is  really  the  an- 
hydride of  chromic  acid,  H2CrO4.  It  is  a  powerful  oxidizing 
agent.  Chromic  acid  is  unknown. 

Chromic  hydroxide,  Cr(OH)3,  is  a  green  solid,  formed  when 
an  alkaline  hydroxide,  carbonate,  or  sulphide  is  added  to  a 
chromic  salt  solution.  It  is  soluble  in  an  excess  of  the  alkali, 
forming  a  chromite  (cf.  §  426).  Chromites  are  derived  from  the 
substance  having  the  formula  HCrO2  ;  this  is  chromic  hydrox- 
ide minus  water.  The  carbonate  and  sulphide  of  chromium 
are  decomposed  by  water,  like  the  corresponding  aluminum 
salts  (cf.  §  427). 

447.  Chromous    and    Chromic    Salts.  —  Chromous 
salts  are  so  readily  oxidized  that  they  are  very  hard  to 
prepare.     In    this  respect   chromium  differs  decidedly 
from  manganese,  the  corresponding  salts  of  which,  the 
manganous    salts,    are    stable    (cf.    §  441).     The    chief 
chromic  salts  are  chromic  chloride  (CrCl3)  and  chrome- 
alum,  K2Cr2(SO4)4.  24  H2O. 

Chromic  chloride  is  obtained  as  a  beautiful,  lavender-colored 


DOUBLE  NATURE  OF  CHROMIUM.  397 

solid  by  passing  chlorine  over  a  mixture  of  chromic  oxide  and 
carbon.  The  carbon  reduces  the  oxide,  and,  at  the  same  time, 
ehlorine  unites  with  the  chromium. 

Chromic  chloride  is  formed  in  solution  by  reducing  a  chrom- 
ate  or  dichromate  with  alcohol,  hydrogen  sulphide,  sulphurous 
acid,  etc.  The  hydrated  salt  is  green. 

2  K2CrO4  +  3  H2S  +  10  HC1  =  4  KC1  -f  2  CrCl3  + 
8  H2O  +  3  S. 

Chrome-alum  is  a  violet-colored,  crystalline  substance, 
formed  as  a  by-product  in  certain  operations  in  which  potas- 
sium dichromate  is  used  as  an  oxidizing  agent.  It  is  analogous 
to  ordinary  alum,  but  contains  chromium  instead  of  aluminum 
(c/-  §  427). 


448.  Double  Nature  of  Chromium.  —  Chromium  is 
not  only  a  metal,  but  also  an  acid-forming  element.  Its 
lower  oxides,  like  those  of  manganese,  form  salts  with 
acids,  and  are,  therefore,  basic ;  but  its  higher  oxides, 
especially  the  trioxide,  CrO3,  are  the  anhydrides  of 
acids. 

The  chromites  formed  by  the  reaction  of  chromic 
hydroxide  with  alkalies  (ef.  §  446)  are  not  important ; 
although  ferrous  chromite,  Fe(CrO2)2,  is  the  chief  chrom- 
ium ore.  The  ehr ornate*  and  the  dickromateg,  however, 
are  the  most  important  chromium  compounds.  Chrom- 
ium is  in  the  same  periodic  group  with  sulphur ;  chromic 
acid  (H2CrO4)  corresponds  with  sulphuric  acid  and 
dichromic  acid,  H2Cr2O7  or  H2O.  2  CrQ3,  with  fuming 
or  disulphuric  acid,  H2S2OT. 


398  MANGANESE  AND    CHROMIUM. 

449.  Chromates  and  Bichromates. —  Potassium  chrotn- 
ate,  K2CrO4,  is  made  by  roasting  ehromite,  Fe(CrO2)2 
or  FeO.  Cr2O3,  with  potassium ,  carbonate  and  quick- 
lime in  the  oxidizing  flame  of  a  reverberatory  furnace 
(fjf.  Fig.  60).  On  a  small  scale  the  ehromite  is  heated 
with  a  mixture  of  potassium  nitrate  and  carbonate. 

To  understand  the  formation  of  the  chruiuate  we  must  think 
of  it  as  made  up  as  follows :  — 

2  K2O  -f  O2Og  +  30=2  K2OO4. 

The  potassium  oxide  is  present  in  the  carbonate,  the  chromic 
oxide  conies  from  the  ehromite,  and  the  oxygen  from  the  oxidiz- 
ing agent  or  the  air.  In  general,  we  change  a  chromic  com- 
pound to  a  chromate  by  oxidizing  it  in  the  presence  of  a  base. 

Potassium  chromate  is  yellow,  like  most  chromates. 
Acids  change  it  to  potassium  diehromate,  K2O2O7. 

Potassium  diehromate  forms  large,  red  crystals,  which 
are  soluble  in  about  ten  parts  of  water  at  the  ordinary 
temperature.  Alkalies  change  it  to  the  chromate. 

To  understand  the  relation  between  the  chromates  and 
dichromates  we  must  look  upon  their  molecules  as  made  up  as 
follows  :  — 

K2Cr,07  =  K20.  2Cr03. 
K2Cr04  =  K20.  Cr03. 

Addition   of   alkalies,  e.  gr.,   potassium    hydroxide,  to   the   di- 
ehromate produces  chromate. 

K20.  2  Cr03  +  2  KOH  (f.  e.,  K2O.  H2O)  =  H2O  -f  2  K2OO4 
(i.  e.,  2  K2O.  CrO3). 

With  acids,  however,  the  reverse  change  takes  place. 

2  K2O.  CrO3  -f  H2SO4  =  K2O.  2  CrO8  -f  K2SO4  -f  H2O. 


OXIDATION   BY  CHBOMATES.  399 

Sodium  dichromate,  Na2O2O7,  is  often  employed  in 
place  of  the  potassium  salt,  owing  to  its  greater  solu- 
bility. 

Both  the  chromates  and  the  dichromates  are  used  in 
dyeing,  in  calico-printing,  and  as  oxidizing  agents.  Po- 
tassium dichromate  is  used  in  photography.  Solutions 
of  chromates  and  dichromates  precipitate  many  metals, 
e.  g.,  lead,  silver,  and  barium,  as  chromates. 

450.  Oxidation  by  Chromates  and   Dichromates.  - 

When  a  chromate  or  a  dichromate  is  used  as  an  oxidiz- 
ing agent,  the  reactions  are  the  reverse  of  those  that 
take  place  when  a  chromate  is  synthesized :  the  chrom- 
ate (or  dichromate)  is  reduced  in  the  presence  of  an 
acid  to  a  chromic  salt. 

Thus  we  think  of  potassium  chromate  as  breaking  down  into 
potassium  oxide,  chromic  oxide,  and  oxygen. 

2  K2CrO4  =  K2O  +  Cr2O3  +  3  O. 

The  oxygen  is  available  for  oxidation ;  while  the  oxides 
unite  with  the  acids,  giving  salts. 

If  the  acid  present  is  sulphuric  acid,  potassium  sulphate 
(K2SO4)  and  chromic  sulphate,  Cr2(SO4)3,  are  formed.  These 
unite  to  produce  chrome-alum,  K2Cr2(SO4)4.  24  H2O. 

If,  however,  hydrochloric  acid  is  present,  we  get  a  mixture 
of  the  chlorides  of  chromium  and  potassium.  If  there  is  no 
other  reducing  agent  present,  the  nascent  oxygen  attacks  the 
hydrochloric  acid,  giving  chlorine  and  water. 

K2Cr207  +  14  HC1  =  2  KC1  +  2  CrCl3  +  7  H2O  -f  3  C12. 


CHAPTER  XXXI. 

LEAD,    TIN,   AND   PLATINUM. 

A.     Lead  (Atomic  Mass,  206.9.) 

451.  Occurrence  and  Preparation  of  Lead. — Lead 
occurs  chiefly  as  galena,  or  galenite,  PbS ;  and  is   ob- 
tained from  it  by  the  following  process  :  — 

The  galena  is  first  roasted  in  a  reverberatory  furnace 
(cf.  Fig.  60).  By  this  operation  part  of  the  ore  is  changed  to 
the  oxide,  PbO,  and  part  to  the  sulphate,  PbSO4,  while  some 
remains  unchanged. 

(1)  2  PbS  -f  3  O2 »  2  PbO  -f  2  SO2. 

(2)  PbS  +  2  O2 >  PbSO4. 

After  the  oxidation  has  gone  far  enough,  the  furnace  doors  are 
closed,  and  the  mixture  is  heated  without  the  admission  of  more 
air.  The  lead  oxide  and  sulphate  then  react  with  the  un- 
changed sulphide  as  follows  :  — 

(1)  PbS  -f  2  PbO »  3  Pb  -f  SO2. 

(2)  PbS04  +  PbS »  2  Pb  -f  2  S02. 

If  the  ores  are  poor,  they  are  often  heated  with  iron. 

PbS -f  Fe  =  FeS -f  Pb. 

If  there  is  enough  silver  to  pay  for  its  extraction,  the  Parkes 
process  is  used  (cf.  §  413). 

452.  Properties  and  Uses.  —  Lead  is  soft,  blue-gray 
metal  having  a  high  luster.     It  is  malleable,  but    not 
very  ductile.     It  melts  at  about  325°  C. 

400 


COMPOUNDS   OF  LEAD.  401 

Although  lead  is  easily  tarnished  in  the  air,  the  cor- 
rosion does  not  penetrate,  as  with  iron.  Ordinary  hard 
waters  act  but  little  upon  lead ;  but  soft  waters  contain- 
ing carbon  dioxide,  organic  matter,  or  much  chloride  or 
nitrate,  attack  it.  Such  waters  should  not  be  carried 
through  lead  pipes  for  household  purposes. 

All  compounds  of  lead  are  poisonous. 

Nitric  acid  acts  readily  upon  lead,  but  hydrochloric 
and  dilute  sulphuric  acids  do  not.  Some  metals,  e.  g., 
zinc,  separate  lead  from  the  solutions  of  its  salts. 

Lead  pipes  are  used  for  conveying  water  and  as 
sheaths  for  the  cables  of  telephone  wires.  In  sheet 
form  the  metal  is  used  to  line  the  "  leaden  chambers  " 
(cf.  §  184),  and  the  sides  and  floors  of  vats  and  tanks 
where  certain  chemical  processes  are  carried  on.  Large 
quantities  of^it  are  made  into  shot,  bullets,  type-metal, 
solder,  pewter,  and  the  plates  of  storage  batteries. 

453.  Compounds  of  Lead.  —  Several  oxides  of  lead 
are  known.  Among  these  are  the  suboxide  (Pb2O),  the 
monoxide  (PbO),  the  dioxide  (PbO2),  and  "  red  lead  "  or 
"  minium,"  which  is  Pb3O4. 

Lead  monoxide,  or  "litharge,"  is  formed  by  heating  lead  in 
a  current  of  air.  At  400°  C.  it  takes  up  more  oxygen,  forming 
"  red  lead,"  Pb3O4.  When  red  lead  is  treated  with  dilute  nitric 
acid,  lead  dioxide  remains  as  a  brown  powder.  Lead  dioxide 
acts  upon  hydrochloric  acid  to  give  chlorine. 

Lead  nitrate,  Pb(NO3)2,  is  made  from  litharge  and 
nitric  acid ;  lead  acetate,  Pb(C2H3O2)2,  from  litharge 


402  LEAD,    TIN,   AND  PLATINUM. 

and  acetic  acid.  Both  are  white,  crystalline  solids. 
Lead  acetate  is  called  "  sugar  of  lead." 

Lead  sulphate  (PbSO4)  and  lead  chromate  (PbCrO4) 
are  insoluble  in  water.  Lead  chromate  is  called 
"  chrome-yellow." 

Lead  chloride,  PbCl2,  is  difficultly  soluble  in  cold 
water,  but  dissolves  in  hot  water. 

Lead  carbonate,  PbCO3,  is  precipitated  from  solutions 
of  lead  salts  by  ammonium  carbonate  solution;  the 
carbonates  of  sodium  and  potassium,  however,  give  a 
basic  carbonate  instead. 

Basic  lead  carbonate  is  made  by  various  methods,  and 
on  a  large  scale ;  it  is  the  pigment  "  white-lead."  It 
forms  a  good  paint,  but  turns  brown  or  black  in  the 
presence  of  hydrogen  sulphide. 

Lead  sulphide,  PbS,  is  precipitated  from.,  the  solution 
of  a  lead  salt  by  soluble  sulphides  and  by  hydrogen  sul- 
phide. 

Plumbites  and  Plumbates.  —  Lead  hydroxide,  Pb(OH)2,  re- 
acts with  alkalies,  giving plumbites,  e.  g.,  K2PbO2  ;  lead  dioxide 
and  alkalies  give  plunibates,  e.  g.,  K2PbO3.  Normal  plumbic 
acid  would  be  H4PbO4  (c/.  silicic  acid,  §  340).  Its  lead  salt 
is  Pb2PbO4,  i.  e.,  red  lead. 

B.     Tin  (Atomic  Mass,  119.) 

454.  Occurrence  and  Preparation  of  Tin.  —  The  only 
mineral  abundant  enough  to  serve  as  a  source  of  tin  is 
cassiterite,  or  tin-stone,  SnO9.  This  occurs  in  Corn- 


COMPOUNDS  OF   TIN.  403 

wall  (England),  in  Australia,  in  the  island  of  Banca, 
and  in  the  Black  Hills. 

Tin  was  known  in  very  early  times.  Cassiterides  was  an 
ancient  name  for  the  Scilly  Islands,  owing  to  the  fact  that  tin- 
stone was  found  there.  Although  tin  has  been  carried  away 
from  Cornwall  since  the  times  of  the  Phoenicians,  the  mines 
there  are  still  producing  it. 

The  metallurgy  of  tin  consists,  first,  in  roasting  the 
ore,  so  as  to  oxidize  and  remove  arsenic  and  sulphur. 
The  tin  oxide  is  then  reduced  with  coal  in  a  furnace, 
the  metal  being  drawn  off  and  cast  into  bars.  These 
bars  of  impure  tin  are  then  slowly  heated  on  a  sloping 
hearth.  The  tin  melts  and  runs  down  the  hearth, 
leaving  the  unmelted  impurities  behind. 

455.  Properties  and  Uses. — Tin  is  a  white    metal 
having  a  brilliant  luster.     It  does  not  lose  its  luster  in 
the  air.     It  is  soft  and  malleable,  and  melts  at  about 
227°  C. 

Tin  reacts  with  hydrochloric  acid,  giving  stannous 
chloride,  SnCl2.  With  concentrated  sulphuric  acid  it 
gives  stannous  sulphate,  SnSO4,  and  sulphurous  acid. 
Nitric  acid  oxidizes  it  to  metastannic  acid  (H2SnO3). 

The  chief  use  of  tin  is  to  coat  sheet-iron  ;  in  this  way  the 
tin-plate  of  commerce  is  formed.  It  is  also  used  to  protect 
other  metals,  e.  (/.,  copper  and  lead  ;  and  in  making  alloys, 
e.  (/.,  soft  solder,  pewter,  bronze,  bell-metal,  etc. 

456.  Compounds  of  Tin.  —  Tin  forms  stannous  and 
gtannw  compounds.     Examples  of  the  former  are  :  stan- 


404  LEAD,    TIN,   AND   PLATINUM. 

nous  chloride  (SnCl2),  stannous  oxide  (SnO),  and  stan- 
nous  sulphide,  SnS.  The  corresponding  stannic  coin- 
pounds  have  the  formulas  SnCl4,  SnO2,  and  SnS2. 
Stannic  acid,  like  metastannic  acid,  is  H2SnO3. 

Stannous  chloride  is  easily  oxidized  to  stannic  chloride,  and 
is,  therefore,  a  good  reducing  agent.  Thus,  it  reduces  mercuric 
chloride  to  mercurous  chloride,  and  even  to  mercury. 

(1)  2  HgCl2  +  SnCl2  =  SnCl4  -f  2  HgCl. 

(2)  2  HgCl  +  SnCl2  =  SnCl,  -f  2  Hg. 

Stannic  chloride  is  a  liquid.  It  is  made  by  heating  tin  in 
chlorine.  A  solution  of  it  is  obtained  by  treating  tin  with 
aqua  regia. 

Stannic  oxide  is  made  when  stannic  acid  is  heated,  and  when 
tin  is  burned  in  the  air.  This  oxide  shows  its  acid  character, 
and  its  analogy  to  carbon  dioxide,  silicon  dioxide,  and  lead  di- 
oxide, by  reacting  with  molten  alkalies  to  form  stannates,  e.  g., 
sodium  stannate,  Ka2SnO3. 

Stannous  sulphide  is  a  brown  powder,  formed  when  tin-foil 
is  heated  with  sulphur,  and  when  hydrogen  sulphide  is  passed 
into  the  solution  of  a  stannous  salt. 

Stannic  sulphide  separates  as  a  yellow  precipitate  when  hy- 
drogen sulphide  is  passed  into  a  stannic  salt  solution. 

Both  sulphides  react  with  alkaline  sulphides,  forming  sul- 
phostannates,  which  are  soluble  in  water  (cf  §§  314  and  318). 

C.     Platinum  (Atomic  Mass,  195). 

457.  Occurrence  and  Preparation.  —  Platinum  is 
found  native  in  a  few  places,  chiefly  in  western  Siberia. 

Native  platinum  is  usually  mixed  with  five  other  rare 
metals,  all  belonging  to  the  eighth  periodic  group.  These  are  : 


CHLORPLATINIC  ACID.  405 

palladium,  ruthenium,  rhodium,  osmium,  and  iridium.     About 
75%  of  the  ore  is  platinum. 

The  ore  is  treated  with  aqua  regia,  which  reacts  with 
the  platinum  and  some  iridium.  The  resulting  chlor- 
platinic  acid,  H2PtCl6  (cf.  bottom  of  page),  is  treated 
with  ammonium  chloride,  producing  a  precipitate  of 
ammonium  Morplatinate,  (NH4)2PtCl6.  When  this  is 
heated  strongly,  metallic  platinum  results.  The  small 
quantity  of  iridium  is  not  removed. 

458.  Properties  and  Uses.  —  Platinum  is  a  grayish- 
white   metal,  over  21  times   as    heavy  as    water.     An- 
nas no  action  upon  it ;  and  the  temperature  of  the  oxy- 
hydrogen  flame  is  needed  to  melt  it. 

Platinum  is  not  attacked  by  the  common  acids ;  but 
it  reacts  with  aqua  regia  and  with  chlorine-  and  bromine- 
water.  Fused  alkalies  also  act  upon  it. 

The  resistance  of  platinum  to  most  chemicals  and  the 
high  temperature  at  which  it  fuses  make  it  very  useful 
in  the  laboratory.  It  is  used  in  the  form  of  foil,  cru- 
cibles, wire,  and  other  utensils.  Large  retorts  of  plati- 
num are  used  in  concentrating  and  distilling  sulphuric 
acid  (cf.  §  185). 

459.  Chlorplatinic   Acid.  —  When  platinum    is  treated 
with  aqua  regia,  the  platinum  chloride  (PtCl4)  formed  unites 
with   hydrochloric    acid,   giving    chlorplatinic    acid,   H2PtCle. 
The  solution  of  this  substance  gives  with  potassium  salts  and 
with  ammonium  salts  precipitates  of  potassium  chlorplatinate, 
K2PtCl6,    and   of   ammonium   chlorplatinate,    (NH4)2PtCl6,   re- 
spectively.    The  corresponding  sodium  salt  is  soluble. 


LABORATORY    DIRECTIONS. 

(For  the  Student.') 

1.  Provide   yourself  with   an    apron  and  a  pair  of   sleeves 
(rubber  is  the  best  material  for  these)  ;   also  with  soap  and 
towel,  and  a  white  cloth  about  a  yard  square.     The  cloth  is  to 
be  used  for  wiping  apparatus. 

2.  Work  by  yourself ;  and  give   your  own  descriptions,  ob- 
servations, and  calculations,  not  those  of  another. 

3.  Record  at  once  all  the  observations  you  make  in  connec- 
tion with  an  experiment.     See    that   your   notes    contain  the 
answer  to  every  question,  direct  or  implied,  that  occurs  in  the 
laboratory  exercise.     Write  neatly  and  distinctly.     If  the  notes 
of  two  experiments  occur  on  the  same  page,  separate  them  by 
at  least  two  centimeters  of  space. 

4.  Have  a  place  for  everything.     Throw  away  nothing  until 
you  are  sure  you  are  through  with  it.     Throw    nothing  but 
liquids  into  the  sink.     Put  other  waste  materials  into  the  proper 
receptacle. 

5.  If  an  experiment  is  unsatisfactory,  repeat  it  until  you  are 
successful ;  but  first  learn  the  probable  cause  of  your  error. 

6.  When  you  enter  the  laboratory,  examine  your  table,  and 
see  that  everything  has  been  left  as  it  should  be  by  the  persons 
who  share  the  table  with  you.     If  anything  is  wrong,  report 
the  fact  at  once  to  the  instructor. 

When  you  leave,  see  that  the  water  and  the  gas  are  turned 
off,  and  that  everything  on  your  table  is  in  good  order. 


LABORATORY  EXERCISES. 


EXPERIMENT   I. 
THE   BUNSEN  BURNER. 

Apparatus.  —  Bunsen  burner,  test  tube,  test-tube  holder  (see 
note  below). 

Materials.  —  Matches,  water. 

a.  Examine  carefully  the  Bunsen  burner  on  your  desk. 
Take  it  apart,  and  draw  a  sketch  of  each  part. 

b.  Put  the  burner  together,  close  the  holes  at  the  base, 
and  connect  with  gas  supply. 

To  light  the  burner,  turn  on  the  gas  and  then  hold  a 
lighted  match  near  the  side  of  the  burner  and  about  one- 
half  a  centimeter  below  its  mouth.  Note  the  character  of  the 
flame ;  is  it  luminous  or  not  ?  Now  open  the  holes  care- 
fully until  therluminous  region  has  just  disappeared.  This 
is  the  "  Bunsen  "  flame.  For  most  work  it  should  be  7  to  10 
centimeters  (3  to  4  inches)  high.  The  holes  of  the  burner 
should  be  open  far  enough  to  prevent  a  deposit  of  soot  upon 
the  object  heated,  but  not  far  enough  to  cause  the  flame  to 
makfe.  a  noise. 

c.  Introduce  quickly  into  the  center  of  the  Bunsen  flame, 
one-half  a  centimeter  above  the  burner,  the  head  end  of  a 
match.     Result  ?     Is  the  gas  in  this  region  burning  ? 

l 


3  LABORATORY  EXERCISES. 

To  heat  an  object  effectively,  place  it  higher  up  in  the 
flame  /  the  best  place  is  just  above  the  apex  of  the  dark, 
inner  cone  of  unburned  gas.  Locate  this  region. 

d.  Put  5  c.c.  water  into  a  test  tube,  and  make  a  note  of 
the  height  of  the  column  of  water  in  centimeters.     When- 
ever you  are  asked  to  take  2,  5,  10,  etc.,  cubic  centimeters 
of  anything,  refer  to  this  experiment,  and  use  the  length  of 
the  column  just  measured  as  your  unit. 

e.  Heat  the  water  in  the  test  tube  to  boiling.     To  do  this 
properly  have  the  outside  of  the  tube  dry ;  hold  the  tube 
in  the  holder,  and  incline  the  tube  at  an  angle  of  about  45° 
to  the  table  top.     Then  introduce  the  bottom  of  the  tube 
into  the  effective   region  (cf.  c)  of  the  flame.     Heat  only 
the  part  of  the  tube  containing  the  liquid /  if  the  flame 
strikes  the  glass  above  the  liquid  level,  the  tube  may  crack. 

Do  not  hold  the  tube  still,  but  move  it  gently  in  the 
flame.  When  boiling  begins,  raise  the  tube  a  little  above 
the  flame,  —  always  keeping  it  inclined,  —  so  that  the  water 
may  not  "  boil  over." 

/.  These  directions  are  general,  and  will  apply  whenever 
you  heat  liquids  in  test  tubes. 

Note.  —  A  very  convenient  test-tube  holder  can  be  made  by 
folding  a  piece  of  writing  paper  twice,  so  as  to  produce  a  strip 
about  1  cm.  wide  and  10  to  15  cm.  long.  This  is  placed  about 
the  tube  like  a  holder.  The  free  ends  are  held  together  close 
to  the  tube. 


CUTTING  AND  BENDING    GLASS   TUBING.  3 

EXPERIMENT   II. 
CUTTING  AND   BENDING   GLASS   TUBING. 

Apparatus. — Bunsen  burner,  "wing-top"  or  illuminating 
gas  burner,  file. 

Materials.  — Piece  of  soft  glass  tubing  more  than  15  cm. 
long. 

a.  Cut  off  a  piece  of  glass  tubing  15  cm.  long.     To  do 
this,  make  on  the  tubing  a  file  mark  in  a  plane  perpendicular 
to  the  length  of  the  tubing ;  grasp  the  tube  in  both  hands, 
and  place  the  thumb  nails  together  opposite  the  scratch. 
By  pushing  gently  with  the  thumbs  and  at  the  same  time 
pulling  with  the  hands  you  will  succeed  in  breaking  the 
tubing  so  that  the  ends  are  fairly  regular. 

b.  Round  off  both  ends  of  the  15  cm.  tube  by  turning 
them  about  in  the  proper  region  of  the  Bunsen  flame  until 
the  edges  become  red  hot.     Let  the  ends  cool. 

c.  Bend  the  15  cm.  tube  at  its  middle  into  the  form  of  a 
right  angle.     For  this  purpose  use  a  flat  Bunsen  flame  — 
produced  by  a  "  wing- top  "  attachment  —  or  a  flat  illumi- 
nating flame.     . 

Take  the  tube  in  both  hands,  one  at  each  end,  and  hold 
its  central  part  lengthwise  with  and  over  the  flat  flame.  At 
the  same  time  twirl  the  tube  between  thumbs  and  fore- 
fingers. Then  lower  the  tube  —  keep  turning  it  —  into  the 
upper  part  of  the  flame,  and  heat  until  you  find  that  the  glass 
is  fairly  soft.  Then  bend  gently  to  a  right  angle. 

d.  If  you  used  the  Bunsen  flame,  anneal  the  glass  at  the 
bend  by  closing  the  holes  of  the  burner  and  allowing  the  hot 


LABORATORY 


glass  to  cool  first  in  the  smoky  flame.  When  the  bend  is 
covered  with  soot,  support  it  so  that  it  will  not  touch  a 
cold  object.  When  the  tube  is  cold,  wipe  off  the  soot. 


EXPERIMENT   III. 
EFFECT   OF  HEAT   UPON   "RED   PRECIPITATE." 

Apparatus.  —  Small  ignition  tube  of  hard  glass,  rubber  con- 
necting tube,  delivery  tube,  pneumatic  trough,  test  tube,  ring 
stand,  clamp. 

Materials. — Pine  splinter,  red  precipitate. 

a.  In  a  small  tube  of  hard  glass  sealed  at  one  end  and 
abou ;  10  cm.  long  —  "  ignition  tube  "  — -place  a  layer  of  red 
precipitate  not  more  than  one-half  a  centimeter  thick. 

In  a  basin  containing  water,  invert  a  test  tube  of  water. 
See  that  no  ah*  bubbles  remain  in  the  test  tube.  Vessels 
for  holding  water  over  which  gases  are  collected  are  called 
"  pneumatic  troughs." 

Attach  to  the  ignition  tube  by  means  of  a  piece  of  rubber 
tubing  a  delivery  tube  long  enough  to  reach  to  the  bottom 
of  the  pneumatic  trough.  Support  the  ignition  and  deliv- 
ery tubes  so  that  the  closed  end  of  the  ignition  tube  is  only 
a  little  lower  than  its  other  end,  and  so  that  the  red  precipi- 
tate may  be  heated  in  the  hot  portion  of  the  Bunsen  flame. 

b.  Begin  to  heat  slowly,  keeping  the  flame  in  motion. 
Note  any  change  in  color  of  the  red  precipitate.    Afterward 
heat  strongly  with  a  steady  flame  until  all  of  the  powder 
disappears.     Collect  over  water  anything  that  escapes  from 
the  delivery  tube  by  allowing  it  to  displace  the  water  of  the 


SOLUTION,   FILTRATION,   AND  EVAPORATION.      5 

test  tube.     When  the  operation  is  over,  remove  the  delivery 
tube  from  the  water  before  removing  the  flame.     Why  ? 

c.  Cover  the  mouth  of  the  test  tube  under  water  with  the 
thumb,  remove  tube  from  water,  invert,  and  introduce  a 
pine  splinter  with  a  spark  on  the  end  of  it.     Result  ?     Is 
the  gas  in  the  test  tube  air  f 

d.  When  the  ignition  tube  is  cool,  invert  it  and  strike  its 
open  end  sharply  against  the  table.     Result?     What  sub- 
stance is  this  ?     On  what  part  of  the  tube  did  it  collect  ? 
Why? 

e.  If  by  a  chemical  change  we  mean  one  in  which  at  least 
one  new  substance  is  formed,  would  you  call  this  a  chemical 
change,  or  not? 


EXPERIMENT  IV. 
SOLUTION,  FILTRATION,  AND  EVAPORATION. 

Apparatus.  —Glass  rod  15  cm.  long  (unfinished),  file,  two 
beakers  of  about  50  c.c.  capacity,  ring  stand,  wire  gauze,  fun- 
nel, funnel  support  (small  ring  of  ring  stand),  evaporating  dish. 

Materials.  —  Coarse  salt,  filter  paper. 

a.  Make  a  glass  stirring  rod   15  cm.  long,  cutting  off  a 
piece  from  a  larger  one,  just  as  in  Experiment  II,  a.  Round 
off  both  ends  in  the  flame. 

b.  Put  into  a  beaker  about  20  c.c.  cold  water,  add  5  grams 
salt,  and  heat  the  beaker  over  the  flame  until  its  contents 
boil.     Before  heating  the  beaker  see  that  it  is  dry  on  the 
outside,  then  place  it  upon  a  wire  gauze  supported  on  the 
ring    stand.     Move  the  flame  about  under  the  ga-uze  until 
the  beaker  has  become  warm  /  then  put  the  burner  under 


6  LABORATORY  EXERCISES. 

the  center  of  the  beaker.  The  height  of  the  gauze  above 
the  burner  should  be  so  great  that  the  bottom  of  the  beaker 
may  be  a  little  above  the  apex  of  the  dark  inner  region  of 
the  flame. 

Note.  —  Always  follow  these  directions  when  you  are  heating 
a  beaker,  an  evaporating  dish,  or  a  flask,  unless  there  is  some 
special  reason  for  not  doing  so.  What  becomes  of  the  salt? 
Of  the  dirt? 

c.  Next,  filter  the  solution.    You  need  a  funnel,  a  support 
(see  above),  a  filter,  the  glass  rod  made  in  a,  and  a  second ' 
beaker. 

Fold  the  circular  filter  twice  in  lines  at  right  angles  to 
each  other.  Press  the  folded  edges  between  thumb  and 
forefinger,  but  not  between  the  nails.  Open  the  filter  so 
that  it  shall  form  an  inverted  cone  which  just  Jits  the  fun- 
nel. One-half  of  the  conical  surface  is  made  up  of  three  of 
the  quarters  into  which  the  paper  was  folded ;  the  remain- 
ing quarter  of  the  paper  makes  up  the  other  half  of  the 
cone. 

d.  Hold  the  filter  in  place  in  the  funnel,  and  wet  it  com- 
pletely ;  it  should  adhere  everywhere  to  the  inner  surface 
of  the  funnel,  and  its  point  should  extend  a  little  into  the 
stem  of  the  funnel. 

Pour  the  salt  solution  down  the  glass  rod  to  the  filter. 

The  glass  rod  should  touch  the  lip  of  the  beaker ;  and 
the  stem  of  the  funnel  should  touch  the  side  of  the  beaker 
beneath  it. 

Always  follow  these  directions  in  filtering  an  insoluble 
solid  from  a  solution. 

e.  Does  anything  remain  on  the  filter  ?     We  call  it  the 
residue.     What  passes  through  is  the  filtrate. 


HYDROGEN.  7 

A  substance  which  remains  mixed  with  a  liquid,  but  not 
dissolved  in  it,  is  said  to  be  "  suspended  in,"  or  "  held  in 
suspension  by  "  the  liquid. 

A  suspended  substance  becomes,  after  filtration,  a  residue. 

f.  Pour  the  filtrate  of  c  into  an  evaporating  dish,  and 
heat  (for  precautions,  cf.  b)  over  the  flame.  Boil  off  the 
water  until  a  solid  begins  to  separate  out ;  then  set  the  dish 
aside  until  it  is  cold,  or  until  the  next  laboratory  period. 
What  is  the  solid  obtained  ? 

Ts  this  separation  of  the  salt  from  the  dirt  a  physical  or 
a  chemical  operation? 


EXPERIMENT  V. 
HYDROGEN. 

Apparatus.  — Generating  flask,  or  bottle  of  250  c.c.  capacity, 
two-holed  stopper,  funnel  tube,  right-angled  tube,  rubber  con- 
nector, delivery  tube,  pneumatic  trough,  squares  of  glass  or  of 
cardboard,  two  or  more  wide-mouth  collecting  bottles  (250  c.c.). 

Materials.  —  Zinc,  dilute  sulphuric  acid  (one  part  by  volume 
of  acid  to  four  volumes  of  water),  pine  splinter,  cupric  sulphate 
solution. 

a.  To  a  250  c.c.  flask  containing  enough  zinc  to  cover  the 
bottom  fit  a  two-holed  stopper.  One  of  the  holes  is  for  a 
funnel  tube  reaching  to  within  one-half  a  centimeter  of  the 
bottom  of  the  flask  when  the  stopper  is  in  place  ;  the  other 
hole  contains  a  bent  tube  attached  by  a  rubber  connector  to 
a  delivery  tube.  The  delivery  tube  reaches  to  a  pneumatic 
trough  containing  two  bottles  filled  with  water  and  inverted. 
The  level  of  the  water  in  the  trough  should  be  about  two 


8  LABORATORY  EXERCISES. 

centimeters  (one  inch)   higher  than  the  months  of  the  in- 
verted bottles  when  the  bottles  are  in  place. 

b.  To  invert  bottles  in  the  trough  without  letting  in  air, 
fill  them  to  overflowing  with  water,.cover  their  mouths  with 
slips    of  glass   or   cardboard,  press   the   latter  against  the 
bottle,  and  invert  quickly  under  water.     Then  remove  the 
covers. 

To  remove  a  bottle  full  of  gas  from  water,  slip  under  the 
mouth  of  the  bottle,  under  water,  a  glass  or  cardboard 
cover,  and  hold  it  in  place  as  before.  Leave  a  filled  bottle 
with  its  mouth  under  water  until  used,  if  possible. 

Whether  a  bottle  of  gas  shall  be  placed  upright  or  in- 
verted upon  the  table  depends  upon  the  specific  gravity  of 
the  gas. 

c.  Caution.  —  Keep  all  flames  at  least  one  meter  (about 
three  feet)  away  from  the  flask  in  which  hydrogen  is 
made. 

See  that  the  stopper  of  the  generating  flask  is  tight,  and 
add  enough  of  the  dilute  sulphuric  acid  to  immerse  the  lower 
end  of  the  funnel  tube. 

Tell  what  takes  place  in  flask,  funnel  tube,  and  pneumatic 
trough.  Explain  each  phenomenon.  If  action  is  not  vigor- 
ous add  a  few  drops  of  copper  sulphate  solution.  Result  ? 
If  evolution  of  gas  ceases  or  becomes  slow  before  you  are 
through,  add  more  acid.  The  gas  produced  is  hydrogen. 

d.  Fill  the  two  bottles  with  the  gas  and  refill  them  after 
using.     Reject  the  first  bottleful    collected   by   turning    it 
mouth  upward.     Why  not  use   it?    Why  turn  it  mouth 
upward  ? 

Keep  the  second  bottle  inverted  and  introduce  into  its 
middle  part  a  burning  pine  splinter  15  to  20  cm.  long.  Hold 


HYDROGEN.  9 

the  splinter  steady  20  to  30  seconds.  Result  ?  Does  the  gas 
burn  ?  Where  ?  Does  the  splinter  continue  to  burn  in  the 
hydrogen?  Is  hydrogen  combustible  or  a  supporter  of 
combustion  ? 

Turn  a  third  bottle  of  the  gas  mouth  upward  one  minute, 
and  repeat  the  test  with  the  burning  splinter.  Results? 
From  the  result  compare  the  specific  gravity  of  hydrogen 
with  that  of  air. 

e.  Place  the  mouth  of  a  fourth  bottle  of  gas  over  the 
mouth  of  an  upright  bottle  of   air.     Hold  the  bottles  to- 
gether and  reverse  their  positions.     After  one  minute  apply 
a  lighted   match   to   the   lower   bottle.     Result  ?     To .  the 
upper.     Jesuit  ?     What  conclusion  as  to  the  diffusibility  of 
hydrogen  ? 

f.  Have  a  fifth  (and  last)  bottle  only  half  full  of  gas ; 
incline  it,  and  then  raise  it  slowly  from  the  water  so  that  air 
displaces  the  remaining  water.     Carry  bottle,  mouth  down, 
to  a  flame.     Result?     Explain  difference  between  this  re- 
sult and  the  combustion  of  hydrogen  free  from  air. 

g.  From  the  experiment  tell  whether  hydrogen  is  very 
soluble  in  water,  or  not. 

h.  Pour  the  liquid  and  the  unused  zinc  from  the  flask  in- 
to a  beaker.  -'  If  the  zinc  has  all  dissolved,  or  if  there  seems 
to  be  enough  acid  to  dissolve  all  of  it,  add  more  zinc. 
Leave  until  action  ceases. 

i.  Examine  the  beaker;  has  anything  separated  from 
solution?  If  so,  re-dissolve  it  by  heating  the  beaker  on  the 
wire  gauze,  and  filter  hot.  (  Care  /) 

Collect  the  filtrate  in  another  beaker  or  an  evaporating 
dish,  and  let  it  stand  some  hours.  Result  ? 

The  substance  you  obtain  is  crystallized  zinc  sulphate. 


10  LABORATORY  EXERCISES. 

What  two  new  products  resulted  from  the  action  of  zino 
and  dilute  sulphuric  acid  ? 


EXPERIMENT   VI. 
EQUIVALENT  OF  MAGNESIUM. 

Apparatus.  —  Balances,  pneumatic  trough,  wide-mouth  bot- 
tle (250  c.c.),  graduated  jar,  glass  or  cardboard  cover. 

Materials. — Magnesium  wire,  dilute  (5%)  sulphuric  acid. 

a:  In  a  pneumatic  trough  containing  water  to  the  depth 
of  about  3  cm.  place  a  piece  of  magnesium  wire  the  exact 
weight  of  which  is  known.  There  should  be  not  more 
than  0.2  gram. 

b.  Get  the  exact  capacity  in  cubic  centimeters  of  a  wide- 
mouth  bottle  \)j  filling  it  with  water  and  pouring  the  water 
into  a  graduated  vessel.     The  bottle  should  hold  at  least 
250  c.c. 

c.  Fill  the  bottle  with  5%  sulphuric  acid,  and  invert  it  in 
the  pneumatic  trough  as  far  from  the  magnesium  as  pos- 
sible.    See  that  the  bottle  is  free  from  air  bubbles. 

Now  slide  the  mouth  of  the  bottle,  under  water,  over  the 
magnesium.  Result  ? 

d.  When  all  the  metal  has  disappeared,  let  the  collected 
gas  cool  to  room  temperature  for  5   minutes.     Then   add 
water  of   room  temperature  to  the  bowl,   if  necessary,  so 
that  the  level  of  water  in  bottle  and  bowl  shall   be   the 
same.     Why  ? 

Protect  the  bottle  from  the  heat  of  the  hand  by  grasping 
it  with  a  towel ;  then  slip  under  its  mouth  a  glass  or  card- 


OXYGEN.  11 

board  cover,  and  invert  quickly ',  so  as  to  lose  none  of  the 
water  in  the  bottle. 

Bring  a  flame  to  the  mouth  of  the  bottle  at  once.  Re- 
sult? 

The  gas  is  hydrogen.  The  other  product  of  the  reaction 
is  magnesium  sulphate ;  it  remains  in  solution. 

e.  Get  the  volume  of  the  water  remaining  in  the  bottle  by 
means  of  the  graduated  vessel.     Then  obtain  by  difference 
the  volume  of  hydrogen. 

To  get  the  weight  of  the  hydrogen  multiply  its  volume 
m  cubic  centimeters  by  the  weight  of  1  c.c.  Get  the 
weight  of  1  c.c.  under  the  conditions  of  the  experiment 
from  the  teacher.  What  is  your  result? 

f.  Solve  the  following  proportion  for  x  :  weight  of  mag- 
nesium :  weight  of  hydrogen  ::  x  :  1.     Result?     x  will  be 
the  equivalent  of  magnesium,  i.  e.,  the  number  of  grams 
of  magnesium  required  to  liberate  1  gram  of  hydrogen  (in 
this  case  from  dilute  sulphuric  acid). 


„      EXPERIMENT  VTI. 
OXYGEN. 

Apparatus.  — Mortar  and  pestle  (?),  test  tubes,  ring  stand 
and  clamp,  one-holed  stopper,  delivery  tube,  pneumatic  trough, 
3  collecting  bottles,  glass  or  cardboard  cover,  deflagration 
spoon. 

Materials. — Powdered  potassium  chlorate  and  manganese 
dioxide,  pine  splinter,  sulphur,  iron  wire  (picture  cord)  at 
least  15  cm.  long. 


12  LABORATORY  EXERCISES. 

a.  On  a  clean  piece    of  writing   paper  mix    carefully  6 
grams  powdered  potassium  chlorate  with  5  grams  powdered 
manganese    dioxide.     If    the  substances   are    not  found  in 
powdered  form  in  the  laboratory,  grind  them  separately,  in 
clean  mortars,  before  mixing. 

b.  Before  you  use  the  whole  mixture,  test  the  quality  of 
a  sample  (1  c.c.)  by  heating  it  gently  in  an  open  test  tube. 
If  there  is  any  evidence  of  violent  combustion,  or  if  large 
sparks  appear,  reject  the  mixture,  and  make  a  fresh  one.     A 
few  small  sparks  indicate  only  traces  of  dust,  etc. 

c.  If  the  mixture  is  satisfactory,  put  it  into  a  test  tube 
supported  by  a  clamp  attached  to  a  ring  stand.     The  test 
tube  is  then  fitted  with  a  one- holed  stopper  and  a  delivery 
tube  reaching  under  water  in  a  pneumatic  trough. 

Have  3  bottles  filled  with  water  and  inverted  in  the 
trough. 

d.  Heat  the  test  tube  gently  from  the  top  of  the  mixture 
downvmrd.     Regulate  the  flame  so  as  to  keep  the  evolution 
of  gas  steady,  but  not  violent.     Keep  the  flame  in  motion, 
so  as  not  to  soften  the  glass. 

When  the  collecting  bottles  are  full,  first  take  the  delivery 
tube  out  of  the  water,  and  then  remove  the  flame.  Why 
this  precaution  ? 

The  gas  is  oxygen. 

e.  Into  one  bottle  of  the  gas  put  a  glowing  splinter  as  in 
Experiment  III,  b.     Result  ?     Gradually  lower  the  splinter 
into  the  bottle  until  combustion  stops.     What  becomes  of 
the  splinter  ?     Of  the  oxygen  ? 

To  the  contents  of  the  bottle  add  5  c.c.  calcium  hy- 
droxide solution  (lime-water),  cover  with  the  hand,  and 
shake  vigorously.  Result? 


KINDLING    TEMPERATURE.  13 

f.  Note  the  odor  of  the  gas  in  the  second  bottle.    Then  put 
into  the  bottle  a  deflagrating  spoon  containing  burning  sul- 
phur.    Light  the  sulphur  by  holding  the  spoon  in  a  flame. 

Have  a  cardboard  cover  with  a  small  hole  for  the  handle 
of  the  deflagrating  spoon,  and  keep  the  bottle  covered  unti] 
combustion  stops.  Results? 

What  becomes  of  the  sulphur  ?  Of  the  oxygen  ?  Note 
the  odor  of  the  gas  now  in  the  bottle.  Does  this  gas  sup- 
port the  combustion  of  a  splinter  ?  Try  it. 

g.  Have  the  third  bottle  of  oxygen  covered  and  set  up- 
right on  the  table.      Draw  aside  the  cover  for  a  moment 
while  you  pour  in  5  c.c.  sand ;  then  replace  the  cover. 

Melt  the  sulphur  left  in  the  deflagrating  spoon,  and  dip 
into  it  one  end  of  a  piece  of  iron  picture  cord.  Light  the 
sulphur  tip,  and  at  once  hold  the  iron  wire  in  the  bottle  of 
oxygen.  Result  ?  Keep  the  wire  in  the  gas  until  action 
ceases.  Describe  the  product.  Why  was  the  iron  tipped 
with  sulphur? 

EXPERIMENT   VIII. 
„  KINDLING  TEMPERATURE. 

Apparatus. — Wire  gauze  at  least  15  cm.  square,  Bunsen 
burner,  tongs. 

Material.  —  Matches. 

a.  Hold  the  wire  gauze,  by  means  of  your  tongs,  7  cm. 
above  the  Bunsen  burner.  Have  the  holes  of  the  burner 
open  as  for  the  Bunsen  flame.  Now  turn  on  the  gas  and 
bring  a  burning  match  from  above  down  to  the  center  of 
the  gauze.  Result  ? 


14  LABORATORY   EXERCISES. 

Why  does  not  the  gas  below  the  gauze  take  fire?  Is 
there  gas  below  the  gauze  ?  Prove  it. 

b.  Let  the  gauze  cool ;  and  then  bring  it  down  upon  the 
Bunsen  flame  until  the  gauze  is  6  to  7  cm.  above  the  top  of 
the  burner.  Result?  Hold  the  gauze  in  place  until  it  be- 
comes red  hot.  Result  ?  Explain. 


EXPERIMENT   IX. 
ACTION  OF  SODIUM  UPON  WATER. 

Apparatus.  — Tongs,  evaporating  dish. 
Materials.  —  Sodium,  water,  blue  and  red  litmus  paper,  solid 
sodium  hydroxide. 

Caution.  —  Do  not  handle  sodium  with  wet  hands,  or 
with  wet  forceps.  Do  not  put  sodium  into  the  waste  jar. 
On  no  account  leave  any  sodium  on  or  about  your  desk  or 
in  your  locker.  Sodium  is  usually  kept  under  kerosene. 

a.  What  is  the  appearance  of  a  freshly  cut  surface  of 
sodium  ?     Is  sodium  hard  or  soft  ?     Heavy  or  light  ? 

b.  Hold  a  piece  of  sodium  having  a  volume  not  greater 
than  8  to  10  c.mm.  at  ami's  length  by  means  of  iron  tongs 
or  forceps,  and  drop  it  upon  water  in  a  small  evaporating 
dish.     Result? 

Apply  a  lighted  match  —  hold  it  at  arm's  length  —  to  the 
sodium  while  it  is  acting  on  the  water.  Result? 

c.  After  action  has  ceased,  wet  your  fingers  with  the  solu- 
tion, and  rub  them  together.     Result?     If  you  get  no  de- 
cided result,  add  a  second  piece  of  sodium  (dry  hands)  of 
the  same  size  as  the  first,  and  repeat. 


WATER    OF  CRYSTALLIZATION.  15 

d.  Test  the  action  of  a  drop  of  the  solution  upon  a  piece 
of  blue  litmus  paper.     Upon  red  litmus  paper.     Results  ? 

e.  Add  a  small  piece  (same  size  as  sodium  used)  of  so- 
dium hydroxide  to  5  c.c.  water.     Results?     Test  solution 
with  the  fingers  and  with  litmus.     Results  ?     Compare  re- 
sults with  those  in  d.     Conclusion. 


EXPERIMENT   X. 
WATER  OF   CRYSTALLIZATION. 

Apparatus.  —  Test  tubes,  iron  saucer  (sand  bath). 
Materials.  —  Crystals  of  zinc  sulphate,  of  potash  alum  (po- 
tassium aluminum  sulphate),  and  of  cupric  sulphate. 

a.  Place  a  few  crystals  of  zinc  sulphate  in  a  dry  test  tube, 
and  warm  gently.     Results  ?     Is  there  evidence  of  water  ? 
Where? 

b.  Repeat  «,  using  a  crystal  of  potash  alum.     Results  ? 

c.  Note  the  taste  of  another  crystal  of  potash  alum ;  then 
heat  it  strongly  in  an  iron  dish  until  no  further  change 
occurs.     Results  ? 

When  the  ignited  alum  is  cold,  taste  it.  Result?  Place 
it  in  5  c.c.  water  in  a  test  tube,  and  boil  carefully  for  five 
minutes.  When  the  water  is  cool,  taste  it.  Result? 

Assuming  that  heat  simply  drove  off  crystal- water  from 
the  alum,  upon  what  does  the  taste  of  crystalline  alum  seem 
to  depend? 

d.  Heat  a  crystal  of  copper  sulphate  (blue  vitriol)  strongly 
in  an  iron  dish.     Result  ?     When  the  residue  is  cold,  add 
a  few  drops  of  water  to  it.     Result  ?     Explain. 


16  LABORATORY  EXERCISES. 

EXPERIMENT   XL 
EFFLORESCENCE. 

Apparatus.  — Evaporating  dish. 

Materials.  — •  Crystallized  sodium  carbonate  and  sodium  sul- 
phate (Glauber's  salt). 

a.  Expose  a  crystal  or  two  of  sodium  carbonate  to  the  air 
for  at  least  twenty- four  hours.     Result  ? 

b.  Carefully  weigh  your  evaporating  dish,  and  then  weigh 
into  it  accurately  about  5  grams  Glauber's  salt  (sodium  sul- 
phate plus  crystal- water).     Let  stand  for  at  least  twenty- 
four  hours,  and  weigh   again.     Result?     What  change  is 
there  in  the  appearance  of  the  substance?     Record  your 

results  thus:  — 

Grams. 

(1)  Weight  of  evaporating  dish  -j-  Glauber's  salt  = 

(2)  Weight  of  evaporating  dish  alone  = 

(3)  .-.  Weight  of  Glauber's  salt  taken  = 

(4)  Weight  of  evaporating  dish  -f-  residue  = 

.*.  gsfin  or  loss  of  water  [subtract  (4)  from  (1)]  = 


EXPERIMENT   XII. 
DELIQUESCENCE. 

Apparatus.  —  Small  beaker,  evaporating  dish  or  watch  glass. 
Materials.  —  Solid    potassium    hydroxide,  granular  calcium 
chloride. 


EFFECT   OF   TEMPERATURE.  17 

a.  In  a  small  beaker  place  a  piece  of  potassium  hydroxide, 
and  leave  it  exposed  to  the  air  at  least  an  hour.     Result  ? 

b.  Weigh  an  evaporating  dish  or  a  watch  glass  carefully, 
and  then  weigh  into  it  accurately  about  5  grams  anhydrous 
calcium  chloride.     Let  stand  at  least  twenty-four  hours,  and 
weigh  again.     Results?     Record  the  weighings  as  in  Ex- 
periment XI,  b. 


EXPERIMENT   XIII. 

EFFECT  OF  TEMPERATURE  ON  SOLUTION. 
CRYSTALLIZATION. 

Apparatus.  —  Beaker  (50  c.c.),  stirring  rod. 
Materials.  —  Potash  alum,  crystallized  cupric  sulphate  (blue 
vitriol). 

a.  Put  20  c.c.  water  into  a  beaker,  add  10  grams  pow- 
dered alum,  and   stir   two  minutes  with  the  stirring  rod. 
Does  all  the  alum  dissolve  ? 

b.  Heat  the  -beaker  carefully  on  the  wire  gauze,  stirring 
the  contents.     Result?     Conclusion. 

c.  Set  the  beaker  with  the  hot  solution  in  cold  water,  and 
stir  rapidly  until  solution  cools.     Result  ? 

d.  Dry  the  outside  of  the  beaker,  and  heat  again  as  hi  b. 
Result  ?     Let  the  solution  stand  undisturbed  until  it  is  cold. 
Result?     Compare  with  c,  and  account  for  the  difference. 

e.  Repeat  a,  5,  c-,  and  4  with  20  c.c.  water  and  15  grams 
powdered  blue  vitriol.     Results  ? 


18  LABORATORY  EXERCISES. 

EXPERIMENT   XIV. 
PRECIPITATION. 

Apparatus.  —  Test  tubes. 

Materials.  —  Solutions  of  lead  nitrate,  potassium  chromate, 
barium  chloride,  and  calcium  sulphate.  Dilute  sulphuric  acid  ; 
alcohol. 

a.  To  5  c.c.  of  lead  nitrate  solution  in  a  test  tube  add 
an  equal  volume  of  potassium  chromate  solution.     Result  ? 
Let  tube  stand  ten  to  fifteen  minutes.     Result?     The  pre- 
cipitate is  lead  chromate. 

b.  Repeat  a,  putting  together  hot  barium  chloride  solu- 
tion and  dilute  sulphuric  acid.     Result  after  ten  to  fifteen 
minutes  ?     The  precipitate  is  barium  sulphate. 

c.  To  2  c.c.  calcium  sulphate  solution  add  an  equal  volume 
of  alcohol.     Result?    The  precipitate  is  calcium  sulphate. 

Note.  —  The  insoluble  solids  formed  in  a  and  b  are  not  the 
only  products  of  these  reactions  ;  the  other  products  are,  how- 
ever, soluble. 

EXPERIMENT   XV. 
CONSTANT   PROPORTIONS. 

Apparatus. — Evaporating  dishes,  beaker,  balances,  watch 
glass. 

Materials.  —  Crystallized  sodium  carbonate,  dilute  hydro- 
chloric acid  (1  volume  concentrated  acid  to  1  volume  water). 

a.  Weigh  your  evaporating  dish  carefully,  and  then  weigh 
into  it  accurately  about  5  grams  sodium  carbonate  crystals. 
Transfer  the  sodium  carbonate  without  loss  to  a  beaker  cov- 


CONSTANT  PROPORTION'S.  19 

ere<l  with  a  watch  glass ;  then  add  the  dilute  hydrochloric 
acid  a  little  at  a  time.  When  adding  acid  draw  the  watch 
glass  a  little  to  one  side ;  at  other  times  let  it  cover  the 
beaker. 

b.  The  effervescence  (foaming)   is  due  to  the  escape  of 
carbon  dioxide  gas.     When  all  the  crystals  have  dissolved, 
add  a  drop  or  two  more  of  the  acid,  to  be  sure  no  sodium 
carbonate  remains ;  then  pour  the  solution  into  the  weighed 
evaporating  dish.     With  5  c.c.  water,  wash  what  has  spat- 
tered on  the  watch  glass  into  the  beaker,  and  with  this 
water  rinse  what  adheres  to  the  beaker  into  the  evaporating 
dish.     Rinse  the  beaker  with  5  c.c.  more  water,  and  add  the 
rinsings  to  the  evaporating  dish. 

c.  Evaporate  the  solution  to  dryness,  on  a  water  bath  or  a 
steam  bath,  if  possible  ;  otherwise,  on  a  wire  gauze.     If  you 
use  wire  gauze  take  great  care  to  avoid  spattering  either  the 
solution  or  the  solid  which  remains  after  the  water   has 
boiled    away.     If   considerable    spattering   begins,   remove 
the  flame  for  a  moment  and  let  the  dish  cool ;  then  apply 
the  flame  again  gently.     Keep  flame  in  constant  motion  at 
the  end  of  the  process. 

When  the  solid  in  the  dish  is  perfectly  dry,  let  the  dish 
cool  to  the  temperature  of  the  room.  Then  weigh  it  accu- 
rately. 

d.  Record  your  results  thus :  — 

Grams. 

Weight  of  evaporating  dish  -f-  sodium  carbonate  = 
Weight  of  evaporating  dish  alone  = 
.•.  Weight  of  sodium  carbonate  used  3= 
Weight  of  evaporating  dish  -f-  sodium  chloride  = 
Weight  of  evaporating  dish  alone  — 
.*.  Weight  of  sodium  chloride  formed  = 


20  LABORATORY  EXERCISES. 

Get  the  simplest  ratio  between  the  amount  of  sodium  car* 
bonate  taken  and  that  of  sodium  chloride  obtained  as  fol- 
lows: Weight  of  sodium  carbonate  :  weight  of  sodium 
chloride  : :  1  :  as.  Calculate  the  value  of  x  to  two  decimal 
places,  x  =  ? 

e.  Repeat  the  preceding  operations,  weighing  out  accu- 
rately about  8  grams  sodium  carbonate  crystals.  If  the 
volume  of  the  solution  is  too  great  to  go  into  the  evapo- 
rating dish  all  at  once,  evaporate  part  of  the  water  and  then 
add  the  remainder  of  the  solution.  Be  sure  to  rinse. 

Calculate  the  ratio  between  sodium  carbonate  and  sodium 
chloride  as  before.  Compare  the  ratios.  Conclusion? 


EXPERIMENT  XVI. 
CHLORINE. 

Caution.  —  Avoid  inhaling  much  chlorine.  If  you  have 
inhaled  it,  smell  ammonia  cautiously.  If  the  gas  gets  into 
the  room,  sprinkle  a  few  drops  of  ammonia  water  upon  your 
table. 

Apparatus.  — 100  c.c.  flask,  ring  stand,  wire  gauze,  one-holed 
stopper,  two  right-angled  tubes  (one  with  long  arm),  rubber 
connector,  collecting  bottle,  test  tubes,  perforated  cardboard 
cover. 

Materials.  —  Manganese  dioxide  (in  lumps),  concentrated 
hydrochloric  acid,  white  paper,  red  cheese  cloth,  litmus  solu- 
tion, indigo  solution,  potassium  chlorate,  ink,  printed  paper. 

a.  Support  a  100  c.c.  flask  on  a  wire  gauze  in  a  ring  stand. 
The  flask  is  provided  with  a  one-holed  stopper  and  a  de- 


CHLORINE.  21 

livery  tube  bent  twice  at  right  angles.  The  double  bend 
is  produced  by  joining  two  right-angle  tubes  by  means  of  a 
rubber  connector.  The  second  right-angle  tube  is  turned 
down;  its  end  should  be  2  to  3  cm.  above  the  table. 

b.  Put  into  the  flask  5  grams  manganese  dioxide  in  small 
lumps  (so-called  crystalline  manganese  dioxide),  add  20  c.c. 
concentrated   hydrochloric    acid,  and   attach   stopper    and 
delivery  tube. 

Warm  the  flask  gently,  and  fill  a  dry  bottle,  turned  mouth 
up,  with  the  resulting  chlorine  gas.  While  the  bottle  is 
being  filled  keep  it  covered  with  a  piece  of  cardboard ;  the 
cardboard  has  a  hole  for  the  delivery  tube.  You  may  know 
when  the  bottle  is  full  by  the  rise  of  chlorine  to  the  top ; 
white  paper  held  behind  the  bottle  will  help  you. 

c.  Stopper  the  bottle  when  it  is  full,  and  fill  two  dry  test 
tubes  with  the  gas.     Then  pass  the  gas  for  five  minutes  into 
15  c.c.  cold  water  in  a  test  tube.     This  gives  chlorine  water. 

When  you  are  through,  disconnect  the  apparatus  at  once, 
and  wash  the  remaining  manganese  dioxide  twice  with 
water. 

d.  What  is  the  color  of  the  gas  ?     Apply  a  lighted  match 
to  a  test  tube  of  it.     Does  the  gas  burn?     Support  com- 
bustion ? 

e.  Put  into  the  bottle  of  the  gas  a  small  piece  (2  cm. 
square)  of  dry  red  cheese  cloth,  a  wet  piece  of  the  same,  a 
piece  of  paper  containing  print,  and  a  paper  with  ink  marks. 
Leave  10  to  15  minutes,  or  until  next  period.     Results? 

/.  Put  5  c.c.  of  the  solution  of  chlorine  made  in  c  upon  1 
sq.  cm.  of  the  colored  cloth  in  a  test  tube.  Upon  paper 
with  both  print  and  ink  marks  on  it.  Results?  What 
seems  to  be  necessary  in  order  that  chlorine  may  bleach  ? 


22  LABORATORY  EXERCISES. 

g.  Into  a  test  tube  of  the  gas  pour  5  c.c.  cold  water,  close 
the  mouth  of  the  test  tube  tightly  with  the  thumb,  and 
shake  vigorously.  Remove  thumb,  under  water.  Result  ? 
Explain. 

h.  To  1  c.c.  dilute  litmus  solution,  add  chlorine  water 
until  you  get  a  decided  change.  Explain  result.  Repeat, 
using  indigo  solution  instead  of  litmus.  Result? 

i.  An  easy  way  to  make  a  solution  of  chlorine  is  to  treat 
about  1  gram  of  potassium  chlorate  with  5  c.c.  concentrated 
hydrochloric  acid.  If  action  is  slow,  warm  gently.  When 
the  effervescence  is  rapid,  add  10  c.c.  cold  water. 


EXPERIMENT   XVII. 
HYDROCHLORIC  ACID. 

Apparatus.  —  Same  as  in  Experiment  XVI. 

Materials. — Sodium  chloride,  concentrated  sulphuric  acid, 
red  and  blue  litmus  paper,  iron  filings,  silver  nitrate  solu- 
tion, ammonium  hydroxide  solution,  dilute  hydrochloric  acid, 
sodium  chloride  solution,  calcium  chloride  solution. 

a.  Into  a  100  c.c.  flask  with  stopper  and  delivery  tube  as 
in  Experiment  XVI,  put  5  to  7  c.c.  water,  and  add  carefully 
20  c.c.  concentrated  sulphuric  acid.     Result? 

Caution.  —  In  diluting  sulphuric  acid  always  pour  the  acid 
into  the  water. 

b.  Cool  the  diluted  acid  by  holding  the  flask  in  a  stream 
of  running  water.     When  the  acid  is  cold,  put  into  it  about 
15  grams  sodium  chloride. 

c.  Attach  the  stopper  and  the  delivery  tube,  and  place 
the  flask  on  the  wire  gauze  of  a  ring  stand.     Warm  care- 


HYDROCHLORIC  ACID.  23 

fully  with  a  small  flame,  and  fill  a  dry  bottle  with  the  gas, 
—  it  is  hydrochloric  acid  gas,  —  as  in  Experiment  XVI,  /;. 

You  may  know  that  the  bottle  is  full  when  white  fumes 
escape  about  the  cardboard  which  covers  the  collecting  bottle. 

Fill,  also,  a  dry  test  tube,  and  stopper  both  vessels. 

d.  Test  the  gas  at  the  end  of  the  delivery  tube  with  strips 
of  moistened  red  and  blue  litmus  paper.     Results  ? 

Blow  your  breath  against  the  stream  of  gas.  Result? 
Kxplain. 

e.  Now  let  the  end  of  the  delivery  tube  come  just  to  the 
surface  of  5  c.c.  water  in  a  test  tube.     Note  the  appear- 
ance of  the  water  below  the  delivery  tube.     While  the  gas 
is  coming  off  regularly,  raise  the  test  tube  so  that  the  end 
of  the  delivery  tube  is  about  2  cm.  below  the  water  level. 
Do  gas  bubbles  pass  through  the  water?     Why?     Lower 
the  test  tube  again  until  the  delivery  tube  is  at  the  surface, 
and  let  the  gas  run  into  the  water  five  minutes.     Is  there 
any  change  in  temperature  ? 

Finally  remove  the  delivery  tube  from  the  water  and 
extinguish  the  flame.  Save  the  liquid. 

Lower  the  wire  gauze  so  as  to  let  the  flask  cool  out  of 
contact  with  the  gauze.. 

f.  Open  the  test  tube  of  gas  under  water.     Result  ?     Ex- 
plain. 

g.  Hold  a  burning  match  in  the  bottle  of  gas.     Does  the 
gas  burn  ?     Does  it  support  combustion  ? 

h.  Test  the  liquid  obtained  in  e  with  red  and  blue  litmus. 
Compare  results  with  those  given  by  the  gas.  Add  a  drop 
of  the  liquid  to  2  c.c.  water,  and  taste  a  drop  held  on  a  stir- 
ring rod.  Result  ? 


24  LABORATORY  EXERCISES. 

i.  Pour  some  of  the  liquid  of  e  upon  1  c.c.  iron  filings  in 
a  test  tube.  Result  ?  Does  the  gas  burn  ? 

The  equation  is :  Fe  -f  2  HC1  ——»  FeCL  -f  ? 
J.  Add  a  few  drops  of    the   liquid    of    e  to  1  c.c.  silver 
nitrate  solution  in  a  test  tube.     Result  ?     The  white  precipi- 
tate is  silver  chloride,  AgCl. 

The  equation  is  : 

AgNO3        +  HC1  =  AgCl  +  ~HNOs. 

silver  nitrate         hydrochloric  acid  nitric  acid 

To  the  precipitate  add  an  excess  of  ammonium  hydroxide 
solution,  close  the  test  tube  with  the  thumb,  and  shake  vig- 
orously. Result  ? 

k.  Repeat  J,  using  sodium  chloride  solution  in  place  of 
hydrochloric  acid.  Result? 

I.  Repeat  j  again  with  calcium  chloride  solution  in  place 
of  the  acid.  Results?  Conclusion? 

m.  Note  the  white  solid  which  separates  when  the  flask 
becomes  cool.  It  is  chiefly  sodium  hydrogen  sulphate, 
NaHS04. 

The  reaction  between  sodium  chloride  and  sulphuric  acid 
takes  place  according  to  the  equation,— 

NaCl  +  H2SO4 »  NaHSO4  -f  HC1. 


EXPERIMENT   XVIII. 
PROPERTIES   OF  ACIDS. 

Apparatus.  —  Stirring  rod,  test  tubes  or  beakers. 
Materials. — Nitric,   sulphuric,  and    tartaric   acids;    litmus 
paper,  iron  tilings. 


PROPERTIES   OF  ACIDS.  25 

a.  Make  a  very  dilute  solution  of  sulphuric  acid  by  add- 
ing three  drops  of  the  concentrated  acid  to  10  c.c.  water  in 
a  clean  vessel.     By  means  of  a  clean  stirring  rod  bring  a 
drop  of  this  solution  to  the  tongue.     What   is   its   taste? 

Note.  —  Whenever  you  taste  a  substance  in  this  way  always 
be  sure  that  it  is  greatly  diluted.  Keep  it  in  the  mouth  long 
enough  to  determine  the  taste  definitely  ;  then  reject  it. 

b.  By   means    of  the    stirring  rod  —  it  must  be   washed 
after  every  test  —  bring  a  drop  of  the  dilute  acid  of  a  upon 
red  and  blue  litmus  papers.     Results?     A  solution  which 
turns  neutral  or  blue  litmus  red  is  said  to  have  an  acid  re- 
action. 

Note. — Litmus  paper  should  not  be  wasted.  One  piece 
will  do  for  many  tests,  if  you  use  only  a,  drop  of  the  liquid  each 
time.  A  new  place  on  the  litmus  paper  must,  of  course,  be 
used  at  every  trial.  To  avoid  mistakes  by  reason  of  substances 
which  may  have  spilled  upon  the  table,  lay  the  litmus  paper 
upon  the  bottom  of  a  clean,  inverted  beaker. 

c.  Try  the  same  experiments  as  in  a  and  b  with  nitric 
acid.     Results?     With   tartaric    acid.     In   the   case  of  the 
tartaric  acid  use  the  solution  obtained  by  heating  a  small 
crystal  with  t)  c.c.  water.     Results? 

d.  From  Experiment  XVII,  *,  tell  what  happens  when 
iron  is  treated  with  hydrochloric  acid.     Gaseous  product  ? 
Write  the  equation  here. 

From  Experiment  V  tell  what  products  are  formed  from 
zinc  and  dilute  sulphuric  acid.  Write  the  equation. 

From  Experiment  VI  tell  what  products  are  formed  from 
magnesium  and  dilute  sulphuric  acid.  If  magnesium  sul- 
phate is  MgSO4,  write  the  equation. 


26  LABORATORY  EXERCISES. 

e.  What  seems  to  be  the  common  gaseous  product  forme< 
when  metals  act  upon  acids  ? 


EXPERIMENT   XIX. 
PROPERTIES   OF   BASES. 

Apparatus.  —  Same  as  in  Experiment  XVI LI. 

Materials.  —  Solid  sodium  hydroxide,  potassium  hydroxide, 
and  calcium  hydroxide,  ammonium  hydroxide  solution,  litmus, 
filter  paper. 

a.  Dissolve  a  small  piece  of  sodium  hydroxide,  NaOH,  in 
10  c.c.  water.     Hub  a  drop  of  the  solution  between  the  fin- 
gers.    Result?     Dilute  3  drops  of  this   with   5    c.c.    water 
and  taste  the  solution,  using  a  stirring  rod.     Result?     Find 
its  effect  upon  blue  and  red  litmus  as  in  Experiment  XVIII, 
b.     Result? 

A  solution  which  turns  sensitive  neutral  or  red  litmus 
blue  is  said  to  have  an  alkaline  reaction. 

b.  Repeat  a,  using  potassium  hydroxide  instead  of  sodium 
hydroxide.     Results  ? 

c.  Add     2     drops    of    ammonium     hydroxide     solution, 
NII4OII,  to  5  c.c.  water.     Note  taste  of  dilute  solution  and 
its  action  on  blue  and  red  litmus. 

d.  Treat    about    1    gram    calcium    hydroxide,    Ca(OII)2, 
or  calcium  oxide,  CaO,  with  10  c.c.  water,  stir  one  minute, 
and  then  filter.     Examine  the  solution  —  it  is  called   lime- 
water —  as   to    taste,  feel,  and    action   upon    blue  and    red 
litmus. 


PROPERTIES   OF  SALTS. 


27 


EXPERIMENT   XX. 
PROPERTIES   OF   SALTS. 

Apparatus.  —  Same  as  in  Experiment  XVIII. 

Materials.  —  Sodium  chloride,  ammonium  nitrate,  potassium 
sulphate,  sodium  acetate,  sodium  carbonate,  disodium  hydro- 
gen phosphate. 

a.  Treat  about  one  cubic  centimeter  of  sodium  chloride, 
N'aCl,  in  a  test  tube  with  5   c.c.  water.     Test  the   solution 
with  blue  and  red  litmus  as  in  Experiment  XVIII,  a  and 
b.     Results  ? 

b.  Repeat  a  with  ammonium  nitrate,  NH4NO8 ;  with  potas- 
sium   sulphate,    K2SO4 ;   with   sodium    acetate,    NaC2Hg()2. 
Results  ? 

c.  Repeat  «,  using  sodium  carbonate,  Na2CO8;  disodium 
hydrogen  phosphate,  Na2HPO4. 

d.  Arrange  in  a  table  the  reactions  of  the  substances  you 
have  examined  with  litmus  in  Experiments  XVII,  XVIII, 
XIX,  and  XX,  thus:  — 


FORMULA 

OF 

SUBSTANCE. 

ACTION  UPON 
BED  LITMUS. 

ACTION  UPON 
BLUE  LITMUS. 

HC1 

etc. 

NaOH 

etc. 

NaCl 
etc. 

• 

28  LABORATORY  EXERCISES. 

What  element  is  found  in  every  acid?  What  two 
elements  are  found  in  every  basic  hydroxide?  What 
element  is  not  present,  usually,  in  the  salts,  i.  e.,  the 
substances  studied  in  this  experiment? 


EXPERIMENT   XXL 
NEUTRALIZATION. 

Apparatus.  —  Evaporating  dish,  stirring  rod,  wire  gauze, 
ring  stand. 

Materials.  —  Litmus  solution  and  paper,  sodium  hydroxide 
solution,  dilute  hydrochloric  and  nitric  acids. 

a.  To  5  c.c.  ten  per  cent  sodium  hydroxide  solution  in  an 
evaporating  dish  add  1  c.c.  litmus  solution ;  then  add  slowly 
dilute  hydrochloric  acid  until  the  litmus  changes  color. 

During  the  addition  of  acid,  stir  constantly  with  a  glass 
stirring  rod.  If  you  get  too  much  acid,  add  sodium  hy- 
droxide by  means  of  the  stirring  rod  until  the  color  just 
becomes  blue  again ;  then  add  a  small  drop  of  very  dilute 
hydrochloric  acid.  With  care  you  can  get  the  litmus  to 
assume  a  color  intermediate  between  the  red  and  the  blue, 
viz. :  a  decided  lavender.  This  is  the  color  of  neutral  lit- 
mus, and  its  formation  shows  that  the  basic  properties  of  the 
sodium  hydroxide  solution  have  been  neutralized  by  the 
hydrochloric  acid. 

b.  Put  a  drop  of  the  solution  upon  red  litmus,  as  in  Ex- 
periment XVIII,  b;  upon  blue  litmus.     Results? 

c..  Evaporate  the  solution  carefully.     At  the  end,  when 


NORMAL  AND  ACID   SALTS.  29 

the  water  is  nearly  all  off  and  spattering  begins,  heat  with 
a  small  name  in  constant  motion. 

d.  Examine  the  product  obtained  in  c,  noting  its  taste, 
solubility  in  water,  and  the  reaction  of  the  solution  with  lit- 
mus.    Results  ? 

The  substance  obtained  is  sodium  chloride,  or  common 
salt.     Complete  the  equation,  NaOH  -\-  HC1  =  ?-(-? 

e.  Repeat  a,  £,  c,  and  d,  using  dilute  nitric  acid  instead  oi 
hydrochloric  acid.     Results? 

If  the   product  has  the    formula  NaNO3,    complete   the 
equation,  — 

-f  HNO3  = 


EXPERIMENT    XXII. 
NORMAL   AND  ACID   SALTS. 

Apparatus.  —  Two  evaporating  dishes,  burette,  test  tube, 
rubber  band,  filter  paper. 

Materials.  — Pure  concentrated  sulphuric  acid,  ten  per  cent 
potassium  hydroxide  solution. 

a.  Put  a  small  rubber  band  evenly  around  a  test  tube  to 
mark  off  5  c.c.  (see  Experiment  I).     Do  not  change  the 
position  of  the  rubber  during  the  experiment. 

b.  Dilute  15  c.c.  pure  concentrated  sulphuric  acid  by  pour- 
ing it  into  35  c.c.  water ;  stir  the  mixture  with  a  glass  rod, 
and  cool  it  as  in  Experiment  XVII,  b. 

Hold  your  marked  test  tube  vertically  and  pour  in  the 
dilute  acid  up  to  the  mark.  See  that  the  upper  edge  of 
the  rubber  is  just  at  the  lower  level  of  the  meniscus,  i.  e.. 


30  LABORATORY  EXERCISES. 

the  curved  surface  of  the  liquid.  Pour  the  5  c.c.  of  acid 
into  an  evaporating  dish,  rinse  the  test  tube  with  5  c.c.  of 
water,  and  add  the  rinsings  to  the  acid  in  the  dish. 

Note  what  part  of  your  evaporating  dish  is  occupied  by 
the  resulting  10  c.c.  of  liquid,  for  comparison  in  e  of  this 
experiment.  Add  to  the  evaporating  dish  1  c.c.  of  litmus 
solution. 

c.  Fill  a  burette  with  ten  per  cent  potassium  hydroxide 
solution.     The  burette  is  best  fitted  with  rubber  and  a  glass 
tip  controlled   by  a  pinch   clamp.     If   a  glass  stopcock  is 
used  see  that  it  is  well  lubricated  with  vaseline.     Support 
the  burette  in  a  clamp,  put  under  it  a  beaker,  and  let  the 
liquid  run  out  until  the  part  of  the  burette  below  the  clamp 
is  filled  with  'liquid.     Return  the  liquid  which  ran  out  to  the 
burette.     Read  the  level  of  the  liquid  exactly  to  tenths  of 
a  cubic  centimeter,  having  your  eye  in  the  same  horizontal 
plane  with  the  bottom  of  the  meniscus.     Record  this  reading. 

d.  Open  the  clamp  of  the  burette  carefully,  and  let  the 
potassium   hydroxide   solution  fall   drop  by  drop  into  the 
evaporating    dish    of    dilute    acid.     Stir    constantly.     Get 
the  solution  exactly  neutral,  or  at  any  rate  have  only  one 
drop  of  alkali  in  excess. 

Read  the  burette  again.     How  much  alkali  was  used? 

e.  Evaporate  the  solution  to  about  1'2  c.c,,  and  let  it  cool 
thoroughly.     Result  ? 

f.  Repeat  #,  c,  and  d  with  twice  the  quantity  of  dilute 
acid,  *.  e.,  10  c.c.,  and  exactly  as  much  potassium  hydroxide 
as  was  used  in  d. 

What  does  the  litmus  now  indicate?  Evaporate  the  re- 
sulting solution  to  5  c.c.,  and  let  it  cool.  Result  ? 

g.  Dry  the  solid  substance  obtained  in  e  between  filter 


NORMAL  AND  ACID  SALTS   CONTINUED.          31 

papers.  What  is  the  general  shape  of  the  crystals?  Heat 
one  in  a  dry  test  tube.  Has  it  crystal- water  ?  Treat  some 
of  the  crystals  with  1  to  2  c.c.  water  in  a  test  tube.  Are 
they  easily  soluble  ?  What  is  the  reaction  of  the  solution 
to  litmus  ?  Its  taste  ? 

h.  Treat  the  crystals  obtained  in  /  as  directed  in  g.     Re 
suits  ?     Are  the  crystals  in  the  two  cases  alike  ? 

How  many  salts  does  sulphuric  acid  form  with  potassium 
hydroxide,  according  to  this  experiment  ? 


EXPERIMENT   XXIII. 
NORMAL   AND   ACID   SALTS   CONTINUED. 

Apparatus.  —  Same  as  in  Experiment  XXII. 
Materials.  —  Concentrated  pure  hydrochloric  acid,  ten  per 
cent  potassium  hydroxide  solution. 

a.  In  the  marked  test  tube  (see  Experiment  XXII,  a) 
measure  out  5  c.c.  concentrated  hydrochloric  acid,  put  this 
acid  into  an  evaporating  dish,  and  rinse  the  tube  with  5  c.c. 
of  water,  as  in  Experiment  XXII,  b.     Add  litmus  solution, 
and  neutralize  with  ten  per  cent  potassium  hydroxide  from 
the  burette.     Note  the  amount  of  alkali  used. 

b.  Evaporate  the    neutral   solution   to  dryness.     Finally 
heat  the  evaporating  dish  until  no  fumes  of  any  kind  come 
off,  and  even  the  crackling  sound  —  decrepitation  —  practi- 
cally ceases. 

Let  the  dish  cool  thoroughly. 

c.  Examine  the  product,  noting  its  solubility  in  water,  the 
taste  of  the  solution,  and  its  reaction  with  litmus, 


32  LABORATORY  EXERCISES. 

d.  Repeat  a,  5,  and  c  with  the  same  amount  of  alkali,  but 
with  twice  the  quantity,  i.  e.,  10  c.c.,  of  hydrochloric  acid. 

Be  sure  to  evaporate  as  directed  in  b. 

e.  Compare  results  with  those  obtained  in  c.     How  many 
salts  do  you  get  hydrochloric  acid  to  form  with  potassium 
hydroxide  ? 


EXPERIMENT   XXIV. 
NITROGEN. 

Apparatus.  — 100  c.c.  flask,  wire  gauze,  ring  stand,  clamp, 
stopper,  delivery  tube,  pneumatic  trough,  collecting  bottle. 

Materials.  —  Sodium  nitrite,  NaNO2 ;  ammonium  chloride, 
NH4C1. 

a.  Support  a  flask  by  means  of  a  clamp  about  its  neck, 
and  place  under  it  the  wire  gauze.     Put  into  the  flask  4 
grams  sodium  nitrite,  3  grams  ammonium  chloride,  and  20 
c.c.  water. 

Attach  the  stopper  and  delivery  tube ;  the  delivery  tube 
extends  to  a  pneumatic  trough  containing  water  and  an 
inverted  collecting  bottle  full  of  water. 

b.  Have  ready  your  evaporating  dish  full  of  cold  water. 
Heat  the  flask  gently  until  a  regular  but  not  too  rapid  stream 
of  gas  escapes.     If  at  any  time  during  the  heating  the  evo- 
lution of  gas  (nitrogen)  becomes  violent,  remove  the  delivery 
tube  from  the  water,  take  away  the  flame  and  wire  gauze, 
and  bring  the  evaporating  dish  of  cold  water  up  over  the 
bottom  of   the  flask.     Let  two   test   tubes    of   gas   escape 
(why  ?)  ;  then  fill  the  bottle  with  it. 


AMMONIA.  33 

c.  Determine  the  odor  and  color  of  the  gas,  also  its  rela- 
tion to  combustion. 

The  equation  representing  its  preparation  is,  in  part,  — 
NH4C1  +  NaN02  =  NaCl  +  ^2  +  H2O. 


EXPERIMENT  XXV. 
AMMONIA. 

Apparatus. —  Mortar  and  pestle,  stirring  rod,  test  tubes,  100 
c.c.  flask,  stopper,  two  right-angle  tubes,  collecting  bottles, 
gauze. 

Materials.  —  Glue,  slaked  lime,  litmus,  hydrochloric  acid, 
ammonium  chloride,  ammonium  nitrate,  ammonium  sulphate, 
sodium  hydroxide  solution,  potassium  hydroxide  solution. 

a.  Mix  in  a  mortar   about   one-half   gram   glue    and    2 
grams    slaked  lime,  and   heat  the   mixture  in  a  test  tube. 
Hold  in  the  mouth  of  the  tube,  without  touching  the  tube, 
a  piece  of  moist  blue  litmus  paper.     Red  litmus  paper.     A 
glass  rod  which  has  been  dipped  into  concentrated  hydro- 
chloric acid.     Results  ?     Note  odor.     What  is  it  ? 

b.  To  about  one-half  gram  ammonium  chloride  in  a  test 
tube  add  2  c.c.  ten  per  cent  sodium  hydroxide  solution,  and 
warm  gently.     Odor?     Effect  of  gas  on  litmus?     On  a  rod 
wet  with  concentrated  hydrochloric  acid  ? 

c.  Repeat  #,  using  ammonium  nitrate  and  sodium  hydrox- 
ide solution.     Results?     Use  ammonium   sulphate  and  ten 
per   cent    potassium    hydroxide    solution.      Results?     The 
gas  formed  in  the  above  cases  is  ammonia,  NH3. 

(1.  In  a  100  c.c.  flask  mix  10  grams  powdered  ammonium 
chloride  and  20  grams  powdered  quicklime.  Odor?  Sup- 


34  LABORATORY  EXERCISES. 

port  the  flask  on  wire  gauze  and  attach  the  stopper  and  a 
delivery  tube  bent  twice  at  right  angles  (see  Experiment 
XVI,  a).  Have  the  second  right-angled  tube  turned  upward. 
On  a  small  ring  fastened  high  up  on  the  ring  stand,  lay  a 
piece  of  cardboard  with  a  small  hole  in  it ;  through  the  hole 
pass  the  delivery  tube,  and  invert  over  the  delivery  tube 
the  dry  receiver  (bottle)  intended  to  collect  the  ammonia. 

e.  Heat  very  yently.     When  the  bottle  is  full  of  gas,  — 
test  this  by  waving  air  from  the  bottle  toward  the  nose,  — 
cover  it  and  place  it  mouth   down   upon   the   table.     Fill 
three  bottles  with  the  gas.  Now  turn  the  end  of  the  delivery 
tube    down,  so  that  it  just  touches  the   surface  of  10  c.c. 
water  in  a  test  tube. 

After  a  minute  raise  the  test  tube  carefully  about  '2 
cm.  Do  the  bubbles  of  ammonia  rise  to  the  surface  of 
the  water?  Why?  Lower  the  test  tube  again  until  the 
delivery  tube  just  touches  the  water,  and  continue  heating 
the  flask  gently  three  minutes.  Remove  the  test  tube ;  and 
then  extinguish  the  flame.  While  the  flask  is  cooling,  place 
between  it  and  the  wire  gauze  a  dry  cloth.  Why? 

f.  Did  you  notice  any  change  in  the  temperature  of  the 
water  of  the  test  tube  ?     Explain.     Save  this  liquid. 

g.  From  the  method  of  collecting  the  gas  compare   its 
specific  gravity  with  that  of  air.     Is  there  any  evidence  of 
water  in  the  generating  flask  ? 

A.  Test  the  gas  in  the  first  receiver  with  litmus  paper  — 
keep  mouth  of  receiver  down.  Test  relation  of  this  gas  to 
combustion.  Results  ? 

Thrust  up  into  -the  receiver  a  glass  rod  which  has  been 
dipped  in  concentrated  nitric  acid.  Results  ?  The  smoke 
is  ammonium  nitrate,  NH4N08.  Write  the  equation. 


AMMONIA.  35 

t.  Place  the  second  bottle  mouth  downward  in  a  pan  of 
water.  Result  ?  Explain. 

j.  Warm  the  bottom  and  sides  of  a  clean,  dry  bottle 
(having  a  mouth  of  the  same  size  as  that  of  the  third  bottle 
of  ammonia)  by  moving  it  quickly  to  and  fro  in  the  Bunsen 
name  ;  put  into  it  five  drops  concentrated  hydrochloric  acid, 
and  place  over  the  bottle  of  hydrochloric  acid  gas  thus 
obtained,  the  bottle  of  ammonia.  Hold  the  mouths  of  the 
bottles  firmly  together  and  reverse  then*  positions,  so  that 
the  ammonia  bottle  is  below  the  other.  Results  ?  What  is 
the  product  ? 

Jc.  Examine  the  solution  of  ammonia  made  in  e.  What 
effect  has  it  upon  litmus  ?  Hold  a  piece  of  moist  red  litmus 
about  2  cm.  above  the  solution.  Result  ?  Explain. 

Put  5  c.c.  of  the  solution  into  a  beaker,  note  the  odor,  and 
let  beaker  stand  for  twenty-four  hours.  Is  the  odor  as 
strong  as  before  ?  Inference  ? 

1.  Put  about  5  c.c.  of  the  ammonia  solution  of  e  into  an 
evaporating  dish,  and  boil  it  gently  for  five  minutes.  Com- 
pare odor  after  boiling  with  that  of  some  of  the  original 
solution. 

m.  Heat  a  small  amount  of  ammonium  chloride  for  some 
time  on  a  piece  of  porcelain  or  on  platinum.  Result  ? 

n.  The  equation  for  the  reaction  which  took  place  in  d 
and  e  is,  — 

OH   ,  NH4C1  __       Cl    ,    NH4OH  /      NH,    ,    H2O\ 

"-  ~~  ° 


Ca(OH)2  +  2  OTI4C1  =  CaCl2  +  2  KH3  +  2  H2O. 
Where  would  you  look  for  the  calcium  chloride  ? 


36  LABORATORY  EXERCISES. 

EXPERIMENT   XXVI. 
NITRIC  ACID. 

Apparatus.  — 100  c.c.  flask,  cork,  stopper,  delivery  tube  (in 
one  piece),  test  tube,  beaker,  wire  gauze,  ring  stand. 

Materials. — -Potassium  nitrate,  concentrated  sulphuric  and 
nitric  acids,  white  silk  thread,  indigo  solution,  ferrous  ammo- 
nium sulphate  or  ferrous  sulphate,  and  copper  nitrate. 

a.  Into  a  100  c.c.  flask  place   about  5  grams  potassium 
nitrate  and  10  c.c.  concentrated  sulphuric  acid.     Attach  the 
stopper  and  the  delivery  tube.     The  delivery  tube  must  be 
in  one  piece  without  rubber  connections. 

Put  the  end  of  the  delivery  tube  into  a  test  tube  resting 
in  a  beaker  of  cold  water.  The  test  tube  will  serve  as  a 
condenser. 

b.  Warm  the  flask  gradually  over  the  wire  gauze.     Re- 
sult ?     Color  of  fumes  ?     Keep  the  end  of  the  delivery  tube 
out  of  the  liquid  which  condenses  in  the  test  tube. 

When  no  more  liquid  distills  over,  remove  the  delivery 
tube.  Only  then  remove  the  flame. 

The  liquid  collected  is  nitric  acid.    Complete  the  equation, 

KNO3  -f  H2SO4 »  KHS04  -f  ? 

What  is  the  color  of  the  acid  ? 

c.  Let   the    flask   cool  thoroughly.     Result?     Name  the 
crystalline  product.     Finally,  add  water  and  pour  the  result- 
ing solution  into  the  sink. 

d.  Add  1  c.c.  of  the  nitric  acid  you  have  made  to  1  c.c.  of 
wrater,  and  test  the  action  of  a  drop  (use  the  glass  rod)  upon 
litmus.     Result  ? 


NITRITES.  37 

Into  your  diluted  acid  put  a  piece  of  white  woolen  yarn, 
and  warm  gently.  Remove  the  yarn.  How  has  it  changed? 

To  1  c.c.  of  your  dilute  acid  add  a  few  drops  of  indigo 
solution.  Result  ? 

e.  What  color  does  your  undiluted  acid  give  to  the  skin  ? 
Will  ammonium  hydroxide  remove  the  stain  ? 

f.  Treat  about  2  c.c.  of  ferrous  ammonium  sulphate  or  of 
ferrous  sulphate  in  a  test  tube  with  15  c.c.  water  and  shake 
vigorously.     Take  5  c.c.  of  this  solution  in  a  test  tube,  add 
two  drops  of  dilute  nitric  acid,  incline  the  tube  at  an  angle 
of  about  forty-five  degrees,  and  pour  about  8  c.c.  concen- 
trated sulphuric  acid  down  the  side  of  the  tube.     Describe 
what  takes   place  where  the   concentrated   acid,  which  is 
below,  meets  the  solution. 

g.  Repeat  /*,  using  a  solution  of  potassium  nitrate  instead 
of  nitric  acid.     Result?     Repeat  again  with  cupric  nitrate 
instead  of  the  acid.     Result? 

h.  If  the  test  just  tried  is  a  general  one  for  all  nitrates, 
how  would  you  proceed  to  test  a  solution  for  the  presence 
of  a  nitrate  or  nitric  acid  ? 


EXPERIMENT   XXVII. 
NITRITES. 

Apparatus.  —  Iron  dish,  stout  iron  wire,  beaker,  test  tubes. 
Materials.  — Potassium  nitrate,  lead,  filter  paper,  dilute  sul- 
phuric acid. 

a.  Melt  together  in  a  shallow  iron  dish  10  grams  potas- 
sium nitrate,  KNO3,  with  about  20  grams  of  lead.     Keep  the 


38  LABORATORY  EXERCISES. 

mixture  at  red  heat,  and  stir  twenty  minutes  with  a  stout 
iron  wire  or  a  nail  held  in  iron  tongs. 

b.  When  the  mass  is  cool  add  2(hc.c.  water,  heat  to  boiling 
for  a  few  minutes,  take  out  the  unused  lead,  and  then  filter. 

The  residue  on  the  filter  is  lead  oxide,  PbO.  Its  color? 
The  filtrate  contains  potassium  nitrite,  KNO2,  and  unchanged 
nitrate.  The  reaction  is  a  reduction  of  potassium  nitrate  by 
lead,  as  is  shown  in  the  equation, 

KN03  +  Pb  =  KNOS  +  PbO. 

c.  Treat  the  solution  of   potassium  nitrite  from  b  with 
dilute  sulphuric  acid.     Result?     Treat  some  potassium  ni- 
trate solution  in  a  test  tube  with  dilute  sulphuric  acid,  and 
compare  results. 

d.  The  brown  gas  is  nitrogen  trioxide,  N2O3,  formed  as 
shown  by  the  equation, 

KN02  +  H2S04 »  KHS04  -f  HN02. 

2  HN0  =  N0        H0. 


EXPERIMENT   XXVIII. 
NITROGEN  TETROXIDE. 

Apparatus. — Test  tubes,  doubly  bent  delivery  tube. 
Material.  —  Powdered  lead  nitrate,  Pb(NO3)2. 

a.  Heat  5  grams  powdered  lead  nitrate  carefully  in  a  test 
tube  (use  a  holder),  keeping  the  tube  in  constant  motion. 
Result  ? 

When  the  tube  is  full  of  gas,  attach  a  stopper  and  r.  tie- 


NITRIC  OXIDE.  39 

livery  tube  with  its  longer  arm  turned  down.     Fill  a  dry 
test  tube  with  the  evolved  gas  by  displacement  of  air. 

b.  Invert  the  test  tube  of  gas  in  a  beaker  of  water,  and 
leave  it  a  few  minutes.     Result  ?     Test  the  residual  gas  with 
a  pine  splinter  having  a  spark  on  the  end  of  it.     Result? 

c.  The  brown  gas  is  nitrogen  dioxide,  NO2,  mixed  with 
nitrogen  tetroxide,  N2O4. 

The  equation  is, 

Pb(NO3)2  =  PbO  +  N20rt  (=  "2  NO,)  4.  O. 

d.  The  lead  oxide  in  the  test  tube  may  nearly  all  be  re- 
moved  by  adding  to  the  cold  tube  dilute  nitric  acid   and 
heating  carefully. 


EXPERIMENT   XXIX. 
NITRIC  OXIDE. 

Apparatus.  —  Generating  flask  (250  c.c..),  two  stoppers  (one 
two-holed  and  one  one-holed),  funnel  tube,  delivery  tubes, 
pneumatic  trough,  wide-mouth  collecting  bottles,  cardboard  or 
glass  covers. 

Materials.  —  Copper  (granulated  or  turnings),  nitric  acid, 
ferrous  sulphate,  splinter  of  pine,  red  phosphorus. 

a.  Into  a  generating  flask  or  bottle  put  enough  granulated 
copper  to  cover  the  bottom.  Attach  stopper  containing 
funnel  tube  and  delivery  tube.  Add  through  the  funnel 
tube  enough  dilute  nitric  acid  to  immerse  the  lower  end  of 
the  tube,  and  then  concentrated  acid,  as  necessary,  to  give 
brisk  action.  The  gas  produced  is  nitric  oxide,  NO.  If 
the  acid  is  too  concentrated,  considerable  nitrogen  tetroxide 
is  produced. 


40  LABORATORY  EXERCISES. 

b.  Fill  a  bottle  over  water  with  the  gas,  and  expose  it  to 
the  air.     Result  ?     From  the  result  tell  why  the  gas  in  the 
generating  flask  was  originally  brown. 

c.  Pass  the  gas  about  three  minutes  into  10  c.c.  concen- 
trated ferrous   sulphate   solution   in  a   test  tube.     Result  ? 
Save  the  solution  for  g. 

cl  Collect  a  second  and  a  third  bottle  full  of  the  gas  over 
water.  Cover  one  of  the  bottles,  remove  it  from  the  water, 
turn  it  mouth  upward,  and  put  into  it  a  lighted  splinter. 
Result? 

e.  Repeat  d  with  the  last  bottle  of  the  gas,  using  a  defla- 
grating spoon  containing   briskly  burning  red  phosphorus, 
instead  of  the  splinter.     Compare  results.     Is  nitric  oxide  a 
supporter  of  combustion,  or  not? 

f.  Attach  to  the  test  tube  of  solution  from  c  a  one-holed 
stopper  and  a  delivery  tube.     Warm  gently  and  collect  the 
evolved  gas  Over  water  in  a  test  tube.     Expose  the  gas  to 
the  air.     Result?     Conclusion? 

The  brown  liquid  obtained  in  c  contains  a  compound  of 
ferrous  sulphate  and  nitric  oxide. 

See  Experiment  XXVI,  /,  where  the  formation  of  a  brown 
ring  of  this  compound  was  used  as  a  test  for  nitric  acid  and  the 
nitrates. 

g.  Pour  25  c.c.  of  the  solution  (its  color  ?)  left  in  the  gen- 
erating flask  into  a  beaker,  add  any  unused  copper,  and  let 
the  beaker  stand  (in  a  gas  chamber  if  possible )  until  all 
action  ceases.  There  should  be  an  excess  of  copper.  Pour 
the  liquid  into  an  evaporating  dish,  and  evaporate  on  a  wire 
gauze  to  about  10  c.c.  Dip  into  the  liquid  a  glass  rod, 
and  see  if  the  liquid  which  sticks  to  the  rod  will  solidify  on 


NITEOUS   OXIDE.  4l 

cooling.    If  so,  let  the  dish  cool ;  if  not,  evaporate  off  about 
2  c.c.  more,  and  try  again.     Result? 

The  substance  obtained  is  cupric  nitrate,  Cu(NO3)2. 

The  equation  is, 

3  Cu  +  8  HN03  =  3  Cu(N03)2  +  4  H2O  +  2  NO. 


EXPERIMENT  XXX. 
NITROUS  OXIDE. 

Apparatus. — Test  tubes,  stopper,  delivery  tube,  pneumatic 
trough,  clamp,  ring  stand,  collecting  bottle. 
Materials. — Ammonium  nitrate,  pine  splinter. 

a.  Into  a  test  tube  provided  with  stopper  and  delivery 
tube  put  about  10  grams  ammonium  nitrate,  and  fasten  the 
test  tube  by  a  clamp  to  the  ring  stand.     The  test  tube 
should  be  inclined  at  an  angle  of  about  forty-five  degrees. 

Invert  a  bottle  of  water  (best  icarm)  in  the  pneumatic 
trough,  but  do  not  put  the  delivery  tube  into  the  water 
until  c. 

b.  Heat  the  test  tube  gently  with  a  moving  flame.  Result  ? 
Warm  more.     Result? 

When  a  steady  stream  of  gas  is  evolved,  hold  over  the 
end  of  the  delivery  tube  a  cold  and  dry  beaker.  What 
collects  in  it  ? 

c.  Now  put  the  end  of  the  delivery  tube  into  the  pneu- 
matic trough,  and  fill  the   collecting  bottle  with  the  gas. 
The  gas  is  nitrous  oxide,  N2O. 

Set  the  bottle  of  gas  mouth  upward  and  covered  upon  the 
table,  and  then  fill  a  test  tube  with  the  gas. 


42  LABORATORY  EXERCISES. 

Note. — Be  sure  to  take  the  delivery  tube  out  of  the    water 
before  you  remove  the  flame. 

d.  To  the  test  tube  of  gas  add  5  c.c.  cold  water,  close  the 
tube  tightly  with  the  thumb,  and  shake  vigorously.     Open 
the  tube  under  water.     Result? 

e.  What  is  the   odor  of  the  gas  in  the    bottle  ?     Insert 
into  it  a  pine  splinter  with  a  glowing  tip.     Result? 

What  gas  resembles  nitrous  oxide  in  its  vigorous  support 
of  combustion? 


EXPERIMENT  XXXI. 
PHYSICAL  PROPERTIES  OF  SULPHUR. 

Apparatus. — Test  tubes,  filter,  funnel,  evaporating  dish, 
beaker. 

Materials.  — -Powdered  sulphur,  carbon  disulphicle. 

a.  Test  the  solubility  of    sulphur  as  follows :    In  a  test 
tube  shake  1  c.c.  powdered  sulphur  with  5  c.c.  water ;  filter, 
and  evaporate  the  nitrate  in  an  evaporating  dish.     Result? 
Conclusion  ? 

What  is  the  odor  of  sulphur  ?     Its  taste  ? 

b.  Treat  not  more  than  1  c.c.  powdered  sulphur  in  a  test 
tube  with  5  c.c.  carbon  disulphide. 

Caution.  —  Carbon  disulphide  is  inflammable.  Do  not 
bring  it  near  a  flame. 

Close  the  test  tube  with  the  thumb,  and  shake  it  thor- 
oughly. Result?  Pour  the  contents  of  the  tube  into  a 
small  beaker,  and  set  this  aside  in  a  gas  chamber  (not  in 
your  cupboard)  until  the  next  laboratory  period.  Result  V 
What  is  the  shape  of  the  larger  crystals  ? 


CHEMICAL   PRO  PER  TIES   OF  SULPHUR.  43 

c.  Fill  a  test  tube  one-third  full  of  sulphur,  hold  it  in- 
clined at  an  angle  of  about  forty-five  degrees,  and  heat  it 
carefully.     Note  the  changes   through  which   the   sulphur 
passes  as  you  raise  its  temperature. 

What  is  the  color  of  the  liquid  formed  by  melting  the 
sulphur?  Is  it  viscous  (thick)  or  limpid  (thin;  easily 
poured)  ?  Pour  a  drop  of  it  into  water.  Color  of  the 
product?  Is  it  hard  or  soft?  Sulphur  melts  at  about 
114°  C. 

d.  What  change  does  the  sulphur  undergo  when  you  heat 
it  further  ?     Tilt  the  test  tube  to  an  almost  horizontal  posi- 
tion from  time  to  time  until  you  find  the  point  at  which  the 
liquid   cannot   be   poured.     Then    continue    heating,    and 
notice  that  the  sulphur  becomes  limpid  again. 

Finally,  heat  the  sulphur  to  boiling.  You  will  know 
that  boiling  is  taking  place  when  you  see  the  dark  brown 
liquid  condensing  upon  the  upper  (cooler)  parts  of  the  tube. 
Sulphur  boils  at  446°  to  448°  C. 

e.  Pour  the  boiling  sulphur  into  a  beaker  of  cold  water. 
Result  ?     Color  of  the  product  ?     Is  it  hard  or  soft  ?     Elas- 
tic or  brittle?     Keep  this  for   several  weeks,  noting  from 
day  to  day  any  changes  that  take  place. 


EXPERIMENT   XXXII. 
CHEMICAL  PROPERTIES   OF   SULPHUR. 

Apparatus. — Deflagrating  spoon,  250  c.c.  bottle,  cardboard, 
test  tube. 

Materials.  —  Sulphur,  blue  litmus  paper,  powdered  iron. 


44  LABORATORY  EXERCISES. 

a.  Put  about  1  c.c.  powdered  sulphur  in  a  deflagrating 
spoon,  heat   it    in    the    flame    until    it    burns    briskly,    and 
then  put  the  spoon  into  a  bottle  of  air.     Keep  the  bottle 
covered  with  cardboard  having  a  hole  in  it  for  the  handle 
of  the  spoon. 

Let  the  sulphur  burn  as  long  as  it  will.  Name  the  prod- 
uct of  the  combustion.  What  is  its  physical  state?  Its 
odor?  Try  its  effect  upon  wet  blue  and  red  litmus  papers. 

b.  Mix  in  a  mortar  5.6  grams   powdered    iron   and    3.2 
grams  powdered  sulphur,  and  put  the  mixture  into  a  test 
tube.     Heat  the  lower  portion  of  the  tube  for  a  moment  in 
the   Bunsen   flame.     Result?     When   action  begins,  with- 
draw the  test  tube  from  the  flame.     Describe  all  that  takes 
place. 

c.  When  the  product  is  cool,  break  the  test  tube,  and  re- 
move the  solid  lump.     Describe  the  product.     It  is  ferrous 
sulphide,  FeS. 

Write  the  equation  for  its  formation.  "Save  the  solid  for 
Experiment  XXXIII. 


EXPERIMENT  XXXIII. 
HYDROGEN   SULPHIDE. 

Apparatus.  — •  Test  tubes,  stopper,  and  delivery  tube. 

Materials.  —  Ferrous  sulphide  from  Experiment  XXXII ; 
dilute  sulphuric  acid  ;  solutions  of  cupric  sulphate,  barium 
chloride,  lead  nitrate,  sodium  hydroxide,  and  litmus. 

Note.  —  Perform  this  experiment  in  a  gas  chamber,  or  where 
there  is  a  good  draught. 


HYDROGEN  SULPHIDE.  45 

a.  Treat  the  lump  of  ferrous  sulphide  made  in  Experi- 
ment XXXII  with  dilute  sulphuric  acid  in  a  test  tube.  Re- 
sult? The  gas  is  hydrogen  sulphide,  H2S.  Attach  a 
stopper  and  delivery  tube,  and  fill  a  test  tube  held  mouth 
upward  with  the  gas.  Apply  a  match  to  the  test  tube. 
Result?  Note  odor  of  the  burning  gas.  Result?  Com- 
plete the  equation, 


b.  Pass  the  hydrogen  sulphide  gas  from  the  generating 
tube  into  5  c.c.  dilute  cupric  sulphate  (CuSO4)  solution  in  a 
test  tube.     Result  ?     Continue  about  one  minute. 

Now  see  that  the  delivery  tube  is  clean,  and  pass  the  gas 
three  or  four  minutes  into  15  c.c.  water  in  a  test  tube. 
Then  wash  out  the  generating  tube  thoroughly. 

c.  Filter  the  test  tube  of  cupric  sulphate  into  which  you 
have  passed  hydrogen  sulphide.     Compare  the  color  of  the 
nitrate    with   that    of    the    cupric    sulphate    taken.     Con- 
clusion ? 

The  black  residue  is  cupric  sulphide,  CuS.  Write  the 
equation  for  the  action  of  hydrogen  sulphide  upon  cupric 
sulphate. 

d.  Add  a  few  drops  of  the  hydrogen  sulphide  solution  to 
2    c.c.   lead    nitrate    solution,   Pb(NO3)2.     Result?     If   the 
insoluble   product   is   lead    sulphide,  PbS,  write  the  equa- 
tion. 

Repeat,  using  cadmium  sulphate  solution  in  place  of  lead 
nitrate.  Result?  If  the  insoluble  product  is  cadmium 
sulphide,  CdS,  write  the  equation. 

e.  Test   the  reaction  of  the  hydrogen  sulphide  solution 
with  red  and  blue  litmus  papers.     Results  ?     Conclusion  ? 


46  LABORATORY  EXERCISES. 

Add  to  the  remainder  of  the  hydrogen  sulphide  solution 
1  c.c.  sodium  hydroxide  solution.  The  solution  now  con- 
tains sodium  sulphide,  Na2S.  Write  the  equation. 

f.  How  would  you  make  ammonium  sulphide  solution, 
(NH4)2S  ?  Write  the  equation. 


EXPERIMENT   XXXIV. 
SULPHUR  DIOXIDE. 

Apparatus.  — 100  c.c.  flask,  ring  stand,  wire  gauze,  stopper 
and  delivery  tube,  2  collecting  bottles,  test  tubes,  beaker,  evapo- 
rating dish. 

Materials.  —  Granulated  copper,  concentrated  sulphuric  acid, 
red  flower,  red  cheese  cloth,  crystals  of  cupric  sulphate,  dilute 
sulphuric  acid,  sodium  hydroxide  solution,  litmus  paper,  con- 
centrated nitric  acid,  potassium  permanganate  and  potassium 
dichromate  solutions. 

a.  In  a  100  c.c.  flask  put  about  5  grams  copper  and  add 
25  c.c.  concentrated  sulphuric  acid.     Support  the  flask  in  a 
ring  stand,  upon  wire  gauze,  and  attach  a  stopper  and  a 
doubly  bent  delivery  tube  reaching  almost  to  the  table. 

b.  Heat  the  flask  carefully.     When  brisk   effervescence 
begins,  moderate  the  heat.     Collect  2  bottles  of  the  gas  as  you 
did  chlorine  in  Experiment  XVI,  b.     Tell  when  each  bottle 
is  full  by  the  odor.     Stopper  the  bottles.     Collect,  also,  a 
test  tube  of  the  gas  and  put  it,  mouth  down,  into  a  beaker  of 
water.     Explain  the  result. 

Wave  a  little  of  the  escaping  gas  toward  the  nose. 
Odor? 


SULPHUE  DIOXIDE.  47 

The  gas  is  sulphur  dioxide,  SO2. 

c.  Put  the  end  of  the  delivery  tube  just  at  the  surface  of 
10  c.c.  water  in  a  test  tube.     When  the  gas  is  coming  off 
freely,  raise  the  test  tube  about  1  cm.     What  evidence  is 
there  that  the  gas  is  dissolving  ?     Lower  the  test  tube  to  its 
former  position  and  keep  it  there  5  minutes.     Then  remove 
the  delivery  tube  from  the  water,  extinguish  the  flame,  and 
let  the  generating  flask  cool  in  position,  out  of  contact  icith 
the  wire  gauze. 

Stopper  the  test  tube  containing  the  solution  of  the  gas, 
and  keep  it. 

d.  Into  one  bottle  of  sulphur  dioxide  gas  put  a  few  petals 
of  some  red  flower,  e.  </.,  a  carnation;  also  a  small  piece  of 
wet,  red  cloth  such  as  you  used  with  chlorine  in  Experiment 
XVI,  d.     Results? 

Test  the  action  of  sulphur  dioxide  upon  blue  litmus  paper. 
Result  ? 

e.  To  the  second  bottle  of  sulphur  dioxide  add  4  drops  of 
concentrated  nitric  acid,  stopper  the  bottle,  and  shake  it. 
Results?     Add  5  c.c  water,  stopper  once  more,  shake,  and 
pour  the  liquid  into  a  test  tube.     Save  this  for  Experiment 
XXXVI,  c. 

f.  Note  the  taste  of  a  drop  of  the  sulphur  dioxide  solution 
made  in  c.     What  is  the  action  of  1  c.c.  of  it  upon   1   c.c. 
potassium  permanganate  solution  (KMn()4)  ?     Repeat  with 
potassium  dichromate  solution  instead  of  potassium  perman- 
ganate.    Result  ? 

Sulphur    dioxide    solution    contains    sulphurous    acid, 
H2S03. 

g.  Neutralize  the  remainder  of  the  sulphurous  acid  in  an 
evaporating  dish  with  10%  sodium  hydroxide  solution  and 


48  LABORATORY  EXERCISES. 

evaporate  to  dryness.     The  resulting  substance  is  sodium 
sulphite,  Na2SO3.     Describe  it. 

Complete  the  equation, 

2  NaOH  -f  H2SO3  =  ?  -f  ? 

h.  Treat  the  sodium  sulphite  in  a  test  tube  with  a  little 
dilute  sulphuric  acid,  and  warm.  Note  the  odor. 

The  sulphite  is  decomposed  by  the  acid  thus:  — 

Na2SO3  +  H2SO4 »  Na2SO4  +  H2SO3  (t.  &,  SO2  -f 

H20). 

i.  When  the  generating  flask  is  cold,  add  to  it  25  c.c.  water, 
shake  carefully,  and  heat  the  flask  cautiously  over  wire 
gauze.  Filter  the  resulting  liquid.  What  is  the  color  of 
the  filtrate?  Concentrate  it  to  about  15  c.c.,  and  let  it  cool. 
Result?  Compare  the  product  with  blue  vitriol,  CuSO4. 
5  H2O.  Complete  the  equation, 

Cu-f  2H2SO4=:2H2O+?  +? 

EXPERIMENT   XXXV. 
SULPHURIC   ACID. 

Apparatus.'- — File  or  blue  paraffin  pencil,  test  tube,  beaker, 
balances.  . 

Materials.  — Concentrated  sulphuric  acid. 

a.  By  means  of  a  file  or  a  blue  paraffin  pencil  mark  off 
about  10  c.c.  on  a  clean,  dry  test  tube,  set  the  tube  in  a  clean 
beaker,  and  get  the  weight  of  both  test  tube  and  beaker 
together.  Now  fill  the  tube  up  to  the  mark  with  concen- 
trated sulphuric  acid,  wipe  off  any  acid  adhering  to  the 
mouth  of  the  test  tube,  and  get  the  weight  of  acid  -j-  beaker 
-|-  test  tube. 


SULPHATES.  49 

Return  the  acid  to  the  bottle,  rinse  the  test  tube,  and  dry 
it  on  the  outside.  Then  fill  the  tube  up  to  the  mark  with 
water,  and  get  the  weight  of  water  -|-  beaker  -j-  test  tube. 

Record  your  results  thus  :  — 

Grams, 

Weight  of  test  tube  -|-  beaker  -}-  sulphuric  acid  = 
Weight  of  test  tube  -J-  beaker  = 
.  •  .  weight  of  sulphuric  acid  taken  = 
height  of  test  tube  -j-  beaker  -j-  water  = 
Weight  of  test  tube  -(-  beaker  = 
.  • .  Weight  of  water  taken  = 

From  the  results  calculate  the  specific  gravity  of  your 
sulphuric  acid. 

b.  Heat  one  drop  —  no  more  —  of  concentrated  sulphuric 
acid  in  an  evaporating  dish  over  wire  gauze.     Result? 

c.  Put  into  a  test  tube  a  splinter  of  wood  and  add  5  c.c. 
concentrated  sulphuric  acid.     Let  stand  15  minutes.     Try 
the  effect  of  a  drop  of  concentrated  sulphuric  acid  on  paper ; 
upon  cotton  cloth.     Wait  for  the  result  if  it  is  not  immedi- 
ate.    Results  ? 

d.  Into  a  small  beaker  put  10  grams  sugar  and   5  c.c. 
water,  and  stir  thoroughly.     Now  add  10  c.c.  concentrated 
sulphuric  acid.     Results  ?     Describe  the  product. 


EXPERIMENT   XXXYI. 
SULPHATES. 

Apparatus.  —  Test  tubes. 

Materials.  —  Concentrated  sulphuric  acid  ;  solutions  of  barium 
chloride,  cupric  sulphate,  sodium  sulphate  ;  dilute  hyhrochloric 
acid  ;  liquid  from  Experiment  XXXIV,  e. 


50  LABORATORY  EXERCISES. 

a.  To  10  c.c.  water  in  a  test  tube  add  1  c.c.  concentrated 
sulphuric  acid.     Result?     Heat  the  diluted  acid  to  boiling, 
and  add  5  c.c.  barium  chloride  solution,  BaCl2.     Result? 

The  precipitated  substance  is  barium  sulphate,  BaSO4.  If 
the  other  product  is  hydrochloric  acid,  write  the  equation. 

Let  the  precipitate  settle,  pour  off  the  supernatant  liquid, 
and  add  10  c.c.  dilute  hydrochloric  acid  to  the  precipitate. 
Does  the  precipitate  dissolve  ? 

b.  Repeat  a,  using  instead  of  dilute  sulphuric  acid  5  c.c. 
of  a  solution  of  cupric  sulphate,  CuSO4.     Results?     Equa- 
tion? 

Repeat  again,  using  5  c.c.  of  sodium  sulphate  solution, 
Na2SO4,  with  5  c.c.  barium  chloride  solution.  Result? 
Equation  ? 

Note.  —  In  general,  if  a  solution  gives  with  barium  chloride 
solution  a  white  precipitate  insoluble  in  hydrochloric  acid,  we 
are  reasonably  sure  that  the  unknown  solution  contains  sul- 
phuric acid  or  a  sulphate. 

c.  Treat  the  liquid  obtained  in  Experiment  XXXIV,  e,  with 
barium  chloride  solution.     Result?     What  effect  did  nitric 
acid  have  upon  the  sulphur  dioxide?     Complete  the  equa- 
tion, 


EXPERIMENT   XXXVII. 
CARBON. 

Apparatus.  —  Tongs,   test   tubes,  iron  dish   with  a  cover, 
beaker. 


CAB  BON  DIOXIDE,   I.  51 

Materials.  — Charcoal  (lumps  and  powder),  graphite  (pencil 
lead),  soft  coal,  hydrogen  sulphide  solution,  litmus  solution., 
brown  sugar,  animal  charcoal. 

a.  Hold  a  piece  of  charcoal  in  the  Bunsen  flame  (use 
tongs)  and  describe  its  combustion.     Repeat  with  graphite 
(pencil  lead)  and  with  soft  coal. 

b.  Fill  an  old  test  tube  one-fourth  full  of  bits  of  wood, 
and  heat.      Results  ?     Bring  a  burning  match  to  the  mouth 
of  the  tube.     Result  ?     Describe  the  other  products.     What 
is  the  residue  ? 

c.  Hold    a   piece    of    wood-charcoal   under    water   in    a 
beaker  for  2  minutes.     What  appears  on  its  surface  ?     Con- 
clusion ? 

d.  Heat  powdered  wood  charcoal  or  animal  charcoal  for  5 
minutes  in  a  covered  iron  dish.    Let  it  cool,  and  add  2  c.c.  of 
it  to  5  c.c.  hydrogen  sulphide  solution.     Shake  thoroughly 
and  filter.      Compare  odor  of  filtrate  with  that  of  the  solu- 
tion taken.      Conclusion  ? 

e.  Boil  5    c.c.  litmus   solution  2   minutes  with  2  c.c.  of 
the  freshly  ignited  charcoal,  and  filter.     Result?     Repeat, 
using  5  c.c.  of  a  solution  of  brown  sugar  with  2  c.c.  fresh 
charcoal.     Result  ?. 


EXPERIMENT   XXXVIII. 
CARBON   DIOXIDE,   I. 

Apparatus.  —  Generating  flask,  stopper  with  thistle  tube  and 
delivery  tube,  pneumatic  trough,  beaker,  test  tubes,  and  col- 
lecting bottles. 


52  LABORATORY  EXERCISES. 

Materials.  —  Marble,  hydrochloric  acid,  litmus,  lime-water. 

a.  Place  in  a  flask  enough  marble  (a  form  of  calcium  car- 
bonate, CaCO3)  to  cover  the  bottom,  add  enough  water  to 
close  lower  end  of  the  thistle  tube,  insert  stopper,  and  add 
concentrated    hydrochloric  acid  through   the    thistle    tube. 
Add  more  acid  when   it   is    needed.     Collect   the  carbon 
dioxide  (CO2)  over  water,  rejecting  the  first  bottle  of  the 
gas.     See,  also,  Experiment  XV. 

b.  Put  into  a  bottle  of  the  gas  wet  litmus  paper  (red  and 
blue)  and  a  burning  match.     Results  ? 

c.  Pour  a  bottle  of  the  gas  into  a  beaker  of  air.     Test 
the   gas    in   the   beaker    with   a  burning   match.     Result? 
Conclusion  as  to  the  specific  gravity  of  the  gas  ? 

d.  Fill  a  test  tube  with  the  gas  by  air  displacement,  add 
5  c.c.  cold  water,  close    tube    securely  with    thumb,  shake 
vigorously,  and  open  under  water.     Result  ?     Conclusion  ? 

e.  Pass  the  gas  into  lime-water,  Ca(OH)2.     Result?     Let 
the  precipitate  settle.     It  is  calcium  carbonate.     Its  forma- 
tion with  lime-water  is  a  test  for  carbon  dioxide  (cf.  Experi- 
ment VII,  e).     Now  pass  a  vigorous  stream  of  the  gas  into 
the  tube  five  minutes.     Result?     Boil  the  contents  of  the 
tube.     Result  ? 

EXPERIMENT   XXXIX. 
CARBON  DIOXIDE,    II. 

Apparatus.  — Beakers,  delivery  tube,  test  tubes. 
Materials.  —  Lime-water,  sodium  bicarbonate,  tartaric  acid. 

a.  Mix  2  c.c.  each  of  sodium  bicarbonate  (NaIICO3)  and 
tartaric  acid  (H2C4H4O6)  in  a  mortar.     Is  a  change  appar- 


REDUCTION  BY  CAEBON.  53 

ent?     Put  half  of  the  powder  into  a  test  tube,  and  add 
water.     Result?     Identify  the  gas. 

b.  Put  the  remainder  of  the  mixture   from   a  in  a  test 
tube,  add  10  c.c.  water,  and,  as  soon  as  you  are  able,  im- 
prison the  gas  by  holding  your  thumb  upon  the  mouth  of 
the  test  tube.     Effect  upon  the  effervescence  ?     Now  remove 
your  thumb.     Effect  ?     Explain. 

c.  Blow  your  breath  through  a  delivery  tube  into  5  c.c. 
lime-water.     Result '?     Conclusion  ? 

d.  Expose  5  c.c.  dear  lime-water  to  the  air  for  several 
hours.     Result?     How  does  carbon  dioxide  get  into  the  air 
(cf.  Experiment  VII,  e)  ? 


EXPERIMENT   XL. 

REDUCTION  BY   CARBON.     EFFECT   OF  HEAT   ON 
CARBONATES. 

Apparatus.  —  Ignition  tube,  delivery  tube,  rubber  connector, 
test  tubes. 

Materials.  —  Lead  monoxide,  powdered  charcoal,  lime-water, 
magnesite. 

a.  Mix  1  c.c.  lead  monoxide,  PbO,  with  one-third  its 
volume  of  powdered  charcoal,  on  smooth  paper.  Into  the 
ignition  tube  put  enough  of  the  mixture  to  make  a  layer  1 
cm.  thick,  support  the  tube  almost  horizontally,  and  attach  a 
delivery  tube  leading  into  5  c.c.  lime-water.  Heat  the  lead 
monoxide  persistently  for  10  minutes,  cool  it,  and  pour  it 
out  on  the  table.  Result  ?  What  gas  was  evolved  ?  Write 
the  equation. 


54  LA  B  OR  A  TOR  Y  EXEE  CISES. 

b.  Fill  the  ignition  tube  one-fifth  full  of  chips  of  mag- 
nesite,  MgCO3,  and  set  it  up  as  in  a.  Heat  persistently. 
What  gas  is  evolved?  What,  then,  does  the  residue  con- 
tain ?  Write  the  equation. 


EXPERIMENT   XLI. 
FLAMES. 

Apparatus.  —  Bunsen  burner  and  tongs. 

Materials.  —  Candle,  piece  of  porcelain,  white  paper. 

a.  Examine  carefully  the  non-luminous  flame.     Sketch  a 
vertical  section  of  it  as  you  see  it.     Make  drawing  4  cm. 
long. 

b.  Do  the  same  with  a  luminous  Bunsen  flame  2  cm.  high. 
Repeat  with  a  candle  flame. 

c.  Press  the  colorless  Bunsen  flame  for  a  moment  upon 
paper  lying  on  your  table.     The  paper  should  not  burn  up. 
Result? 

Hold  a  piece  of  glass  tubing  about  1  dm.  long  at  an  angle 
of  45°,  with  the  lower  end  inside  the  central  part  of  the 
non-luminous  flame,  and  apply  a  lighted  match  to  the  other 
end.  Result?  What  do  these  experiments  show  as  to  the 
inner  region  of  the  flame  ? 

d.  Hold  a  piece  of  porcelain  (broken  evaporating  dish)  by 
means   of   tongs   in  the   luminous    flame.     Result?     What 
substance  is  in  excess  here  ?     Now  hold  the  porcelain  iir  the 
colorless  flame  for  some  time.     Result  ?     What  is  in  excess 
in  this  flame  ? 


WEIGHT   OF  A   LITER    OF    OXYGEN. 


55 


EXPERIMENT   XLII. 
WEIGHT   OF   A   LITER   OF  OXYGEN. 

Apparatus.  —  A  strong  test  tube,  two  one-liter  bottles,  bent 
glass  tubes,  pinch-clamp,  graduated  cylinder,  balances. 

Materials.  —  Powdered,  chemically  pure  potassium  chlorate 
and  manganese  dioxide,  dried  at  120°  C.  Water  at  room 
temperature. 

a.  Set  up  the  apparatus  shown  in  Fig.  64.     A  is  a  test 
tube  having  a  strong  flare,  so  that  it  can  be  slipped  tightly 
over  a  rubber  stopper. 

B  is  a  liter  bottle  fitted 

with  a  two-hole  rubber 

stopper.      The    longer 

tube  reaches  almost  to 

the  bottom  of  J?,  and 

is  connected  at  F  with 

a  rubber  tube  reaching 

to   the  bottom   of  the 

bottle  C.     The  rubber 

tube  may  be  closed  by 

the     pinch- clamp     F. 

Almost  fill   the   bottle 

£  with  water  having  the  temperature  of  the  room,  and  then 

fill  the  long  tube  and  the  rubber  tube  by  sucking  at  Z>,  A 

being  removed.     Then  close  the  pinch- clamp. 

b.  Into  the  test  tube  A  put  about  5  c.c  of  a  mixture  of 
equal  parts  of  powdered,  chemically  pure  potassium  chlor- 
ate and  manganese  dioxide ;  they  must  have  been  dried  at 


B 


D 


FIG.  64. 


56  LABORATORY  EXERCISES. 

120°  C.  for  at  least  an  hour.     Get  the  weight  of  test  tube 
and  mixture  accurately  on  the  balances,  and  record  it. 

c.  See  that  the  stopper  is  pressed, securely  into  the  mouth 
of  ./?,  and  then  slip  A  carefully,  but  tightly,  over  its  stopper. 
Now  put  about  50  c.c.  water  into  C,  raise  C  so  that  the  water 
in  £  and  G  are  at  the  same  level,  open  the  pinch- clamp  one 
minute,  and  then  close  it.     Then  put  (J  down  on  the  table. 
Take  the  rubber  tube  carefully  out  of  C  and  get  the  volume 
of  the  water  in  C  /  then  pour  the  water  back  into    C,  and 
put  the  rubber  tube  in  place.     ISTow  open  the  pinch- clamp, 
and  hang  it  upon  the  glass  tube  at  E.     Do  not  allow  the 
lower  end  of  the  rubber  tube  to  get  above  the  surface  of 
the  water  in  C.     Why? 

d.  Heat  the  mixture  in  A  gently,  beginning  at  the  upper 
part  of  the  mixture.     The  evolved  gas  forces  water  from  B 
into  C.     When  C  is  about  half  full,  stop  heating,  and  let  A 
cool  to  room  temperature.     Then  raise  15  or  (7,  as  necessary, 
to  make  the  water  levels  in  both  the  same  (be  sure  to  keep 
the  lower  end  of  the  rubber  tube  under  water),  close  the 
rubber  tube  with  the  pinch- clamp,  and  get  the  volume  of 
the  water  in  C.    This,  minus  the  original  volume,  equals  the 
volume  of  gas  collected  in  7?. 

Find  the  barometric  height,  subtract  from  it  the  cor- 
rection for  the  pressure  of  water  vapor  (see  Appendix),  and 
find,  also,  the  temperature  of  the  gas.  Finally,  weigh  A  and 
its  contents  accurately. 

e.  Record  your  results  thus :  —  Grams. 

Weight  of  test  tube  -|-  contents  at  first  = 
Weight  of  test  tube  -j-  contents  afterward  = 
.*.  Weight  of  oxygen  = 

Volume  of  oxygen  at      °  C.  and      mm.  =  c.c. 

.'.  Volume  of  oxygen  at  0°  C.  and  760  mm.  =  c.c. 

Weight  of  oxygen  obtained  :  weight  of  a  liter  at  0°  C.  and 
760  mm.  ::  volume  (at  0°  C.  and  760  mm.)  :  1,000  c.c. 


BROMINE.  57 

EXPERIMENT   XLIIL 
BROMINE. 

Apparatus.  —  Beaker,  100  c.c.  flask,  test  tubes. 

Materials. — Potassium  bromide,  powdered  manganese  diox- 
ide, dilute  sulphuric  acid,  litmus  paper,  calico,  carbon  disul- 
phide,  chlorine-water,  sodium  hydroxide. 

Caution.  —  If  possible,  work  in  a  gas-chamber  or  hood. 

a.  Into  a  flask  put  an  eighth  of  a  test  tube  of  potassium 
bromide  (KBr)  crystals,  half  as  much  powdered  manganese 
dioxide,  and  half  a  test  tube  of  dilute  sulphuric  acid.     Sup- 
port the  flask  over  wire  gauze,  and  attach  the  cork  stopper 
and  a  doubly  bent  delivery  tube  reaching  into  a  test  tube 
three-fourths  full  of  cold  water.     The  delivery  tube  must 
be  without  rubber  connections.     The  test  tube  should  rest 
in  a  beaker  of  water. 

b.  Warm  the  flask  carefully  until  a  dark  brown  distillate 
passes  over.     Is  it  heavier  or  lighter  than  water  ?     Do  not 
inhale  the  vapor,  and  do  not  get  liquid  bromine  on  your 
hands. 

When  no  more  bromine  comes  over,  remove  first  the  de- 
livery tube  and  then  the  flame.  The  light-brown  solution 
is  "  bromine- water." 

c.  Wave  the  air   from  the  test  tube  toward   the   nose. 
Odor  of  bromine  ?     Pour  off  as  much  bromine- water  as  pos- 
sible without  pouring  out  the  bromine,  and  add  more  water 
to  the  bromine.     Pour  a  few  drops  of  the  bromine- water 
upon  litmus  paper  and  upon  colored  calico.     Results  ? 

d.  To  3  c.c.  water  in  a  test  tube  add  1  c.c.  carbon  disul- 


58  LABORATORY  EXERCISES. 

phide,  close  tube  with  thumb,  and  shake  vigorously.  Re- 
sults? Where  is  the  carbon  disulphide?  Now  add  £  c.c. 
bromine- water  and  shake  again.  Result  to  the  color  of  the 
water  ?  To  that  of  the  carbon  disulphide  ?  This  effect  on 
the  carbon  disulphide  is  a  test  for  free  bromine. 

e.  To  5  c.c.  of  potassium  bromide  solution  add  1  c.c.  car- 
bon disulphide  and  shake.     Result  ?     Now  add  two  or  three 
drops  of  chlorine-water  (made  as  in  Experiment  XVI,  f), 
close  the  tube,  and  shake  it  as  before.     Results  ?     Action  of 
chlorine  on  potassium  bromide  ?     Equation  ? 

f.  To  the  liquid  bromine  in  the  test  tube  add  sodium  hy- 
droxide   solution,    a    c.c.    at   a   time,    shaking    thoroughly. 
(Do  not  close  the  tube  with  the  thumb !)     Result  ?     The 
equation  is, — 

2  NaOH  +  2  Br »  NaBr  -f  NaBrO  +  H2O. 

sodium 
hypobromite 


EXPERIMENT    XLIV. 
IODINE   AND   HYDRIODIC  ACID. 

Apparatus.  —  Test  tubes,  beaker,  flask. 

Materials.  — Potassium  iodide,  manganese  dioxide,  sulphuric 
acid,  iodine,  carbon  disulphide,  chlorine-  and  bromine-water, 
starch,  alcohol,  hydrogen  sulphide,  silver  nitrate  solution,  so- 
dium carbonate,  litmus. 

a.  Powder  potassium  iodide  (KI),  mix  1  c.c.  of  it  with 
a  c.c.  of  manganese  dioxide,  add  2  c.c.  water  and  then 
1  c.c.  concentrated  sulphuric  acid.  Result?  When  the 


IODINE  AND  HYDRIODIC  ACID.  59 

action  slackens,  warm  the  tube  gently,  and  then  let  it  cool. 
Describe  what  you  find  in  the  tube  ?  It  is  iodine.  Com- 
pare its  preparation  with  that  of  chlorine  and  bromine. 

b.  Warm  a  crystal  of  iodine  (gently,  not  to  boiling)  with 
10  c.c.  water  for  a  few  seconds.     Does  the  iodine  dissolve 
readily?     Cool  the  water  and  add  3  c.c.  of  it  to  1  c.c.  car- 
bon disulphide.     Shake  the  closed  tube.     Result  ?     This  is 
a  test  for  free  iodine.     Save  the  iodine  solution. 

c.  Shake  5  c.c.  potassium  iodide  solution  with  1  c.c.  car- 
bon  disulphide.     Result?     Add  a  drop  of   chlorine- water 
and  shake  again.     Result  ?     What  effect  has  chlorine  upon 
potassium  iodide  ?     Repeat,  using  bromine- water  instead  of 
chlorine- water.     Write  both  equations. 

d.  Make  a  starch  solution  as  follows :  Mix  2  c.c.  powdered 
starch  with  5  c.c.  cold  water,  and  pour  the  emulsion  into  30 
c.c.  boiling  water.     Boil  for  a  minute  or  two,  and  then  cool. 
To  3  c.c.  of  the  solution  add  a  drop  of  the  iodine  solution 
of  b,  shaking.     Result  ? 

To  3  c.c.  of  the  starch  solution  add  one  drop  of  a  potas- 
sium iodide  solution  and  then  one  drop  of  chlorine-  or 
bromine- water.  Result  ? 

e.  Heat  a -crystal  or  two  of  iodine  in  a  dry,  inclined  test 
tube.     Result?     Let   cool.     Result?     Effect   of   iodine  on 
the  skin  ?     On  wood  and  paper  ? 

/.  To  the  iodine  of  e  add  5  c.c.  ethyl  alcohol,  C2H5OH. 
In  which  is  iodine  more  soluble,  water  or  alcohol?  An 
alcoholic  solution  is  often  called  a  tincture. 

g.  To  one-half  a  c.c.  of  powdered  iodine  in  a  flask  add  20 
c.c.  water  and  then  pass  in  hydrogen  sulphide  (gas-chamber !) 
until  the  iodine  disappears.  Results?  Boil  the  solution 
gently  two  minutes,  and  filter  it.  Identify  the  precipitate 


60  LABORATORY   EXERCISES. 

by  igniting  a  little  on  a  piece  of  porcelain.  Odor  ?  Test 
the  filtrate  with  red  and  blue  litmus.  Results?  Add  a 
drop  of  it  to  1  c.c.  silver  nitrate  solution.  Result?  Add 
some  to  1  c.c.  solid  sodium  carbonate.  Result  ?  What  sub- 
stances are  formed  from  hydrogen  sulphide  and  iodine  ? 
Equation  ? 


EXPERIMENT   XLV. 
COMPARISON  OF  THE   HALOGEN  ACIDS. 

Materials.  —  Potassium  chloride,  bromide,  and  iodide;  con- 
centrated sulphuric  acid,  litmus. 

a.  Three  test  tubes    have   small   amounts  of   potassium 
chloride,  bromide,  and  iodide,  respectively ;  treat  each  with 
a  few  drops  of  concentrated  sulphuric  acid.     Results  ?     Blow 
your  breath  across  the  mouth  of  each  tube.     Result  ?     Test 
the  gas  of  each  with  blue  litmus.     Result  ?     Note  carefully 
the  odor  of  each  gas.     What  odors  beside  that  of  the  acid 
do  you  get  in  the  tube  of  potassium  iodide?     Heat  this 
tube.     Result  ? 

b.  Which  tube  gives  a  colorless  gas  ?     What  colors  the 
seas  in  each  of  the  two  other  cases  ?     From  the  amount  of 

O 

coloration,  tell  which  of  the   three   halogen   acids   is  most 
easily  decomposed  into  its  elements.     Which  least. 


HYDROGEN  PEROXIDE.  61 

EXPERIMENT   XLVI. 
HYDROGEN  PEROXIDE. 

Materials. — Hydrochloric  acid,  barium  peroxide,  starch  so- 
lution, potassium  iodide  solution,  manganese  dioxide,  potassium 
permanganate,  ether,  potassium  dichromate  solution,  splinter. 

a.  To  25  c.c.  water  add  5  c.c.  concentrated  hydrochloric 
acid  and  then  3  grams  powdered  barium  peroxide,  BaO2,  a 
little  at  a  time,  stirring.     Filter  the  solution ;  it  should  con- 
tain Jiydrogen  peroxide,  H2O2. 

b.  To  5  c.c.  starch  solution  add  a  drop  of  potassium  iodide 
solution,  and   then   a  few  drops   of  the   hydrogen  peroxide 
solution.     Result  ? 

c.  To  3  c.c,  of  the  solution  of  a  add   3  c.c.  ether.     Do 
they  mix?     Is  the  ether  above  or  below?     Now  add  one 
drop  of   potassium    dichromate    solution.     Close   tube   and 
shake  gently.     Result  ? 

d.  To  5  c.c.  of  the  hydrogen  peroxide  solution  add  1  c.c. 
powdered  manganese  dioxide.     Result  ?     Test  gas  with  a 
glowing  splinter.     Result? 

e.  To  three  crystals  of  potassium  permanganate  in  a  test 
tube  add  2  c.c.  water  and  then  5  c.c.  of  the  hydrogen  perox- 
ide solution.     Result  ?    Test  with  glowing  splinter.     Result  ? 


62  LABORATORY  EXERCISES. 

EXPERIMENT   XLVII. 
PHOSPHORUS   AND   PHOSPHORIC  ACID. 

Apparatus.- — Test  tubes,  small  ignition  tube,  tongs,  evapo- 
rating dish,  file. 

Materials.  —  lied  and  yellow  phosphorus,  carbon  disulphide, 
filter  paper,  phosphoric  acid,  ammonium  hydroxide  ;  silver 
nitrate,  disodium  hydrogen  phosphate,  magnesium  sulphate, 
ammonium  chloride,  and  calcium  chloride  solutions. 

Caution. — Ordinary,  yellow  phosphorus  must  be  handled 
only  with  tongs,  never  with  fingers  !  It  must  be  kept  and  cut 
under  water.  No  pieces  of  it  must  get  into  your  locker  ;  and 
every  dish  that  has  contained  phosphorus  must  be  heated,  so 
that  the  phosphorus  may  be  completely  burned. 

Do  not  bring  carbon  diftulphide  near  a  flame  I 

a.  Put  half  a  c.c.  of  red  phosphorus  into  a  test  tube,  and 
add  3  c.c.  carbon  disulphide.     Result?     Filter,  and  let  the 
carbon  disulphide  evaporate,  without  heating,  in  a  hood,  or 
where  its  vapor  will  not  get  near  a  flame.     Result  ?     Was 
any  phosphorus  dissolved  ? 

To  3  c.c.  carbon  disulphide  add  a  piece  of  yellow  phos- 
phorus not  larger  than  a  grain  of  wheat.  Shake  carefully 
a  few  minutes.  Result  ?  Pour  the  solution,  every  drop  of 
it,  upon  a  piece  of  filter  paper  laid  flat  on  a  ring  of  the  ring 
stand.  Let  the  carbon  disulphide  evaporate  without  heating 
it.  Result  ?  Rinse  the  test  tube  before  putting  it  away. 

b.  Into  a  small  ignition  tube  put  a  layer  of  red  phosphorus 
not  more  than  5  mm.  thick,  hold  tube  horizontal  (tongs),  and 
gently  heat  end  containing  the  phosphorus.     What  collects 


ARSENIC.  63 

on  the  cold  part  of  the  tube  ?  When  the  tube  is  cold,  make 
a  file-mark  just  below  the  deposit,  and  break  the  tube.  Rub 
the  deposit  with  a  match  stick.  Result?  Conclusion? 
Finally,  heat  both  tubes  red  hot,  so  as  to  burn  up  all  the 
phosphorus.  Throw  the  pieces  into  iron  or  crockery  jars. 

c.  To  5  c.c.  water  add  1  c.c.  concentrated  (ortho)  phos- 
phoric acid,  neutralize  in  an  evaporating  dish  (use  litmus) 
with   ammonia,   and    add    silver   nitrate   solution.     Result? 
The  precipitate  is  silver  orthophosphate,  Ag3PO4.     Write 
the  two  equations. 

Dissolve  2  c.c.  powdered  sodium  hydrogen  phosphate  in 
10  c.c.  water.  To  half  of  the  solution  add  calcium  chloride 
solution.  Result  ?  The  product  is  secondary  calcium  phos- 
phate, CaHPO4.  Equation  ? 

d.  To  5  c.c.  magnesium  sulphate  solution  add  1  c.c.  am- 
monia-water and   1   c.c.   ammonium  chloride  solution,  and 
then   the    disodium   hydrogen   phosphate   solution   from   c. 
Result  ?     The  product  is  magnesium  ammonium  phosphate^ 
NH4MgP04.  6H20. 

(1)  :Na2HPO4  +  NH4OH »  Na2^H4PO4  -f  H2O. 

(2)  ^a2NH4PO4  -f-  MgS04 »  MgNH^PO^  -f  Na2SO4. 


EXPERIMENT   XL VIII. 
ARSENIC. 

Apparatus.  —  Small  ignition  tube,  tongs,  test  tube,  beaker. 

Materials.  —Arsenic  trioxide  (powdered),  charcoal,  hydro- 
chloric acid,  hydrogen  sulphide,  ammonium  sulphide,  sodium 
hydroxide  solution. 


64  LABORATORY  EXERCISES. 

a.  Into  a  small  ignition  tube  put  powdered  arsenic  triox- 
ide, As2O3,  to  the  depth  of  5  mm.    Hold  the  tube  horizontal 
and  at  the  side  of  the  flame,  so  as^to  heat  only  the  end  con- 
taining the  powder.     What  happens?     Now  slip  into  the 
tube,  almost  to  the  arsenic  trioxide,  a  piece  of  charcoal  about 
2  cm.  long.     Heat  the  charcoal  red  hot  (have  tube  hori- 
zontal), and  then  incline  the  tube  slightly  so  as  to  heat  the 
arsenic  trioxide  while  keeping  the  charcoal  red  hot.     lie- 
suit  ?     Effect  of  charcoal  upon  the  oxide  ?    Equation  ?     How 
does  the  oxide  come  into  contact  with  the  charcoal  ?     Sub- 
lime the  arsenic  obtained. 

b.  Heat  half  a  c.c.  of  arsenic  trioxide  with  8  c.c.  dilute 
hydrochloric  acid   to  gentle   boiling.      Result?     Equation? 
Pour  off  from  any  undissolved  material,  and  pass  in  hydro- 
gen sulphide  for  a  minute.     Result?     If  visible  product  is 
arsenic  trisulphide,  As2S3  (its  color?),  write  equation.     Let 
settle,  pour  off  supernatant  liquid,  and  add  5  c.c.  ammonium 
sulphide  to  residue,  shaking.    Result  ?     (CAUTION.  —  Do  not 
get  ammonium  sulphide  on  your  hands  !  )      The  product 
now  formed  is  ammonium  sulpharsenite,  (NH4)3AsS3;   it  is 
soluble.     Treat  solution  with  an   excess    of    dilute   hydro- 
chloric acid  in  a  beaker.     Result? 

c.  Treat  half  a  c.c.  of  arsenic  trioxide  with  sodium  hy 
droxide  solution.     Warm  carefully.     Result  ?     The  solution 
contains  sodium  arsenite,  Na3AsO3.     From  b  and  c  would 
you  say  arsenic  trioxide  has  acid,  or  basic,  properties  ? 


ANTIMONY.  65 

EXPERIMENT  XLIX. 
ANTIMONY. 

Apparatus. — Mortar  and  pestle,  funnel,  ignition  tube. 

Materials.  —  Antimony,  concentrated  nitric  and  hydrochloric 
acids,  hydrogen  sulphide,  ammonium  sulphide,  antimony  triox- 
ide. 

a.  What  is  the  color  of  metallic  antimony  ?    Is  it  heavy 
or  light?     Powder  a  small  piece,  and  treat  part  of  it  in  a 
test  tube  with  concentrated  nitric  acid.     Results  ? 

b.  Treat  the  remainder  of  the  powdered  antimony  of  a 
with  3  c.c.  concentrated  hydrochloric  acid  and  1  c.c.  concen- 
trated nitric  acid.     Warm  to  start  the  action,  if  necessary. 
The  solution  contains  antimony  chloride,  SbCl3.     Let  action 
continue  for  ten  minutes ;  then  add  15  c.c.  water.     Filter,  if 
necessary,  and  pass  in  hydrogen  sulphide.     If  there  is  no 
action,  dilute  still  more.     Result?     If  the  product  has  the 
formula   Sb2S3,  write   the   equation.     Treat   the  antimony 
sulphide   as    you    did   arsenic    trisulphide   in    Experiment 
XL VIII,  b. 

c.  Dissolve  half  a  c.c.  of  tartar  emetic  in  5  c.c.  water,  add 
a  drop  of  hydrochloric  acid,  and  pass  in  hydrogen  sulphide. 
Result  ?     Conclusion  ? 

d.  Heat  antimony  trioxide  (Sb2O3)  in  an  ignition  tube  with 
charcoal,  as  you  did  arsenic  trioxide.     Results? 


66  LABORATORY  EXERCISES. 

EXPERIMENT   L. 
BISMUTH. 

Apparatus.  — Mortar  and  pestle,  beaker,  test  tubes. 
Materials. — Bismuth,  concentrated  nitric  and  hydrochloric 
acids,  bismuth  nitrate  crystals,  hydrogen  sulphide. 

a.  What  is  the  color  of  bismuth  ?     Is  the  metal  heavy  or 
light  ?     Malleable  or  brittle  (test  with  the  pestle)  ?     Treat  a 
bit  with  concentrated  nitric  acid.     Result  ?     Products  ? 

b.  To  half  a  c.c.  of  bismuth  nitrate  crystals,  Bi(NO3)3, 
add  5  c.c.  water,  and  shake.     Result?     If  the  product  has 
the  formula  BiONO3,  write  the   equation.     Now  add  hy- 
drochloric acid  (concentrated)  a  drop  at  a  time,  heating  to 
boiling  after  each  drop.     Result?     Use  the  least  possible 
amount  of  acid. 

c.  Put  half  of  the  solution  from  b  into  a  beaker,  and  add 
much  water.     Result  ?     Compare  with  first  part  of  b. 

d.  To  the  remainder  of  the  acidified  solution  of  bismuth 
nitrate  from  b  add  hydrogen  sulphide.     Result  ?     The  visible 
product  is  bismuth  sulphide,  Bi2S3.     Write  the  equation. 


EXPERIMENT   LI. 
BORAX  AND  BORIC  ACID. 

Apparatus.  — Platinum  wire  sealed  into  glass  rod,  test  tubes, 
beaker. 

Materials. — -Borax,  potassium  dichromate,  manganese  diox- 
ide, hydrochloric  acid,  and  sodium  carbonate  (solid). 


BORAX  AND    BORIC  ACID.  6? 

a.  Borax  Bead.     Make  a  loop  2  mm.  in  diameter  on  the 
end  of  a  platinum  wire  sealed  into  a  piece  of  glass  tubing. 
Heat  the  loop  to  redness,  and  dip  it  into  powdered  borax, 
Na2B4O7. 10  H2O.     Heat  the  adhering  borax  just  within  the 
outer  edge  of  the  Bunsen  flame,  at  the  place  where  the  flame 
is  widest.     This  is  the  fusing  zone  of  the  flame.     What 
happens  first?     Heat  until  the  borax  melts  to  a  transparent 
glass.     If  there  is  not  enough  borax  to  fill  the  loop,  add 
more,  and  heat  again.     This  glassy  lump  is  called  the  borax 
bead. 

b.  Touch  the  hot  bead  to  a  tiny  speck  (less  than  half  as 
large  as  a  pin's  head)  of  potassium  dichromate,  K2O2O7, 
and  heat  at  the  top  of  the  flame  until  the  dichromate  is  com- 
pletely absorbed  by  the  bead.     Color?     Remove  the  bead 
by  plunging  it  while  hot  into  water,  and  wipe  it  off  the  wire. 

c.  Make  a  new  bead,  and  touch  it  to  a  speck  of   man- 
ganese dioxide.     Heat  first  in  the  oxidizing  flame  of  the 
burner,  i.  e.,  just  above  the  visible  tip  of  the  flame.     Color? 
Now  heat  it  in  the  reducing  region,  i.  e.,  just  above  the  tip 
of  the  bright  blue  interior  cone.     Heat  it  there  persistently 
for  five  minutes,  examining  it  from  time  to  time.     Result? 
Heat  it  again  in  the  oxidizing  flame.     Result  ? 

d.  Boric  Acid. — Dissolve  5  grams  powdered  borax  in  10 
c.c.  hot  water,  and  add   10  c.c.   concentrated  hydrochloric 
acid.     Set  aside  until  next  laboratory  period.     Result  ?     The 
product  is  boric  acid,  H3BO3.     Filter  off  the  crystals,  wash 
them  on  the  filter  with  a  little  cold  water,  and  dry  them  on 
fresh  filter  paper. 

e.  Dissolve  the  crystals  of  boric  acid  in  hot  water,  and  add 
the  solution  to  a  lump  of  sodium  carbonate.     Result  ? 


C8  LABORATORY  EXERCISES. 

EXPERIMENT   LII. 
IONIZATION. 

Materials.  —  Solutions  of  silver  nitrate,  potassium  chloride, 
potassium  chlorate,  ammonia,  and  potassium  ferrocyanide  ;  fer- 
rous sulphate  crystals. 

a.  Take  2  c.c.  silver  nitrate  solution  in  each  of  two  test 
tubes.     To  one  tube  add  a  few  drops  of  potassium  chloride 
solution.     Name  the  precipitate  from  Experiment  XVII,  -j 
to  /.     Equation? 

b.  To  the  second  tube  add  potassium  chlorate  solution. 
Result?     Is  the  same  substance  precipitated  as  in  a? 

c.  Powder  a  crystal  of  ferrous  sulphate,  FeSO4,  and  shake 
it  with  5  c.c.  water.     Pour  off  the  solution,  and  add  to  it  a 
c.c.  of  ammonium  hydroxide  solution.     Result  ?     If  the  pre- 
cipitate is  ferrous  hydroxide,  Fe(OH)2,  write  the  equation. 

d.  To     3    c.c.      of     potassium     ferrocyanide     solution, 
K4Fe(CN)6,  add  ammonium  hydroxide.     Result?     Is   fer- 
rous hydroxide  precipitated,  as  before  ? 

e.  Refer  to  Experiment  XXXIX,  a.     Why  does  the  ac- 
tion between  sodium  bicarbonate  and  tartaric  acid  take  place 
only  when  water  is  present? 


EXPERIMENT   LIIL 
HYDROLYSIS  AND  MASS  ACTION. 

Materials. — Antimony  trichloride    (crystalline   or  melted), 
hydrochloric  acid. 


SODIUM  COMPOUNDS.  69 

a.  To  a  small  amount  (half  a  c.c.)  of  antimony  chloride, 
SbCl3,  add  5  c.c.  water,  and  shake.     Result?     The  visible 
product  is  essentially  antimony  oxychloride,  SbOCl,  i.  e., 

OH 
SbOH  minus  water.     Write   the  equation.     Compare  the 

Cl 
result  with  Experiment  L,  #,  where  bismuth  nitrate  was  used. 

b.  To  the  precipitate  add  concentrated  hydrochloric  acid, 
a  drop  at  a  time,  warming  after   each  drop.     Result?     If 
the  solution  contains  antimony  chloride,  SbCl3,  write  the 
equation. 

c.  Add  the  solution  obtained  in  b  to  50  c.c.  water.     Re- 
sult?    Add  concentrated  hydrochloric  acid  again.     Result? 

d.  Compare   the    equations   of  a  and  b.     Write  one  of 
them,  using,  instead  of  the  equality  sign,  the  double  arrow 
c       *.     In  which  direction  does  the  reaction  go  chiefly  when 
an  excess  of  water  is  used?     When  an  excess  of  acid  is 
used? 


EXPERIMENT   LIV. 
SODIUM   COMPOUNDS. 

Apparatus.  —  Test  tubes,  stopper  and  delivery  tube,  magni- 
fying glass,  platinum  wire,  watch  glass  or  glass  slip. 

Materials. — Sodium  bicarbonate,  lime-water,  sodium  car- 
bonate (solid  and  in  solution),  sodium  chloride,  calcium  chlor- 
ide, barium  chloride,  sodium  nitrate  and  sulphate,  hydrochloric 
acid. 

a.  Refer  to  Experiment  IX  for  the  properties  of  sodium 
and  its  action  on  water. 


70  LABOEATOBY  EXERCISES. 

b.  Heat  2  c.c.  powdered  sodium  bicarbonate  carefully  in  a 
test  tube  having  a  delivery  tube  that  passes  into  lime-water. 
Result?     Is  there  any  other  volatile  product?     When  no 
more  gas  is  evolved  (do  not  melt  the  test  tube),  let  the 
product  in  the  tube  cool,  and  then  add  2  c.c.  cold  water. 
Note  the  temperature  effect.     Compare  with  this  the  action 
of  anhydrous  sodium  carbonate  upon  water.     What  are  the 
products  formed  by  heating  sodium  bicarbonate  ?     Equation  ? 

c.  Heat  2  c.c.  sodium  chloride  with  5  c.c.  water  in  a  test 
tube;    filter;    and  let  some  of   the  filtrate   evaporate   com- 
pletely on   a   glass    slip    or    a  watch    glass.     Examine    the 
crystals  with  a  magnifying  glass,  if  possible.     Their  shape  ? 

d.  Dissolve  a  small  piece  of  calcium,  chloride,  CaCl2,  in 
5  c.c.  water,  and  add  sodium  carbonate  solution.     Result? 
Repeat,  using  barium  chloride  instead  of  calcium  chloride. 
Result?     Write  the  equations. 

e.  Dip  a  platinum  wire  with  a  glass  holder  (cf.  Experi- 
ment LI,  a)  into  5  c.c.  concentrated  hydrochloric  acid  in  a 
test  tube,  and  then  heat  the  wire  in  the  Bunsen  flame  until 
the  flame  remains  colorless.     If  necessary,  dip  the  wire  more 
than  once.     Now  wet  the  clean  wire  with  the  acid,  dip  it 
into  powdered  sodium  chloride,  and  heat  it.     Effect  on  the 
flame? 

f.  Clean  the  wire  and  repeat  e,  using  sodium  nitrate  in- 
stead of  sodium  chloride.     Repeat  again  with  sodium  sul- 
phate.    What  color  do  sodium  salts  give  to  the  flame? 


POTASSIUM  COMPOUNDS.  71 

EXPERIMENT   LV. 
POTASSIUM   COMPOUNDS. 

Apparatus.  —  Watch  glass,  iron  dish,  test  tubes,  beaker  or 
evaporating  dish,  platinum  wire,  copper  wire. 

Materials.  —  Potassium  chloride,  sodium  nitrate,  sulphur, 
barium  chloride  solution,  potassium  hydrogen  tartrate,  lime- 
water,  dilute  sulphuric  acid,  concentrated  hydrochloric  acid, 
potassium  nitrate,  and  potassium  sulphate. 

a.  Heat  8  grams  of  potassium,  chloride  and  10  grams  of 
sodium  nitrate  with  20  grams  of  water  until  there  is  com- 
plete solution,  and  boil  off  half  of  the  water  over  the  wire 
gauze.  Result?  Let  the  precipitate  settle  and  pour  the 
solution  into  a  test  tube.  Wash  the  residue  with  5  c.c.  cold 
water,  and  then  dissolve  it  in  the  smallest  possible  amount 
of  hot  water.  Pour  a  few  drops  of  the  solution  in  a  watch 
glass  and  set  aside.  Result?  Compare  the  crystals  with 
those  obtained  in  Experiment  LIV,  c.  Conclusion  ? 

What  happens  in  the  test  tube  containing  the  original 
solution  ?  The  visible  product  is  potassium  nitrate,  KNO8. 
.  b.  Mix  3  c.c.  powdered  potassium  nitrate  on  a  clean  piece 
of  paper  with  1  c.c.  powdered  sulphur,  and  pour  the  mixture, 
at  arm's  length,  upon  a  hot  iron  dish  (use  no  wire  gauze). 
Result?  Let  the  product  cool,  boil  it  with  10  c.c.  water  in 
a  test  tube,  and  add  to  5  c.c.  of  it  barium  chloride  solution. 
Result  ?  See  Experiment  XXXYI,  b.  What  is  the  product 
of  the  deflagration  of  potassium  nitrate  and  sulphur  ? 

c.  Heat  an  iron  dish  red  hot,  and  pour  upon  it  3  c.c.  pow- 
dered potassium  hydrogen  tartrate,  KHC4H4O6  (cream  of 


72  LABORATORY  EXERCISES. 

tartar).  Results?  Color  of  residue ?  Heat  it  five  minutes 
longer  at  red  heat,  pressing  the  mass  down  with  a  glass  rod 
occasionally.  When  the  dish  is  coql,  treat  the  residue  in  a 
test  tube  with  dilute  sulphuric  acid.  After  all  evolution  of 
gas  ceases,  identify  the  gas  by  placing  in  the  mouth  of  the 
tube  a  looped  copper  wire  holding  a  drop  of  lime-water. 
What  remains  undissolved?  What  substance  would  you 
find  in  plant  ashes  if  the  plants  contained  potassium  salts  of 
organic  acids  ? 

d.  Clean  a  platinum  wire  as  in  Experiment  LIY,  e  ;  dip 
it  into  strong  hydrochloric  acid,  and  then  into  powdered 
potassium  chloride,  and  heat  it  in  the  flame.  Result  ?  Re- 
peat, using  potassium  nitrate  instead  of  the  chloride.  Use 
the  sulphate.  Results?  What  color  do  potassium  com- 
pounds give  to  the  flame  ? 


EXPERIMENT   LVI. 
•SOLUBILITY  OF  POTASSIUM   CHLORIDE. 

Apparatus.  —  Steam  bath,  water  bath,  or  wire  gauze  ;  evap- 
orating dish,  balances. 

Materials.  —  Powdered  potassium  chloride,  distilled  water. 

a.  Make  a  saturated  solution  of  potassium  chloride  by 
shaking  12  grams  of  the  powdered  substance  in  a  clean  flask 
with  25  c.c.  distilled  water  at  the  temperature  of  the  room. 
Continue  shaking  every  little  while  for  fifteen  minutes. 
Record  the  temperature  of  the  solution,  and  then  weigh  out 
accurately  into  your  evaporating  dish  about  20  grams  of  the 
solution.  Now  evaporate  (see  b)  the  water  until  the  residual 


AMMONIUM  AMALGAM.  73 

potassium  chloride  is  perfectly  dry,  and  get  its  weight. 
From  the  results  calculate  how  much  potassium  chloride 
will  dissolve  in  100  grams  of  water  at  the  room  temperature. 

b.  If  possible,  evaporate  the  solution  of  a  on  a  steam  or 
water  bath.     If  this  is  impossible,  evaporate  slowly  and  care- 
fully on  wire  gauze,  so  as  to  avoid  any  loss  by  spattering. 

c.  Record  your  results  thus :  — 

Grams. 

Weight  of  evaporating  dish  -j-  water  -|-  KC1  = 
Weight  of  evaporating  dish  alone  = 
.'.  Weight  of  water  -f  KC1  = 
Weight  of  evaporating  dish  -J-  KC1  = 
Weight  of  evaporating  dish  alone  = 
.'.  Weight  of  KC1  = 
.'.  Weight  of  water  found  :  weight  KC1  found   ::  100  grams 

:  x. 


EXPERIMENT  LVII. 

AMMONIUM   AMALGAM.     DISTINCTIONS  BETWEEN  THE 
ALKALI   METALS. 

Materials^ — Ammonium  chloride,  sodium  amalgam,  sodium 
and  potassium  chlorides,  tartaric  acid,  two  unknown  substances. 

a.  Dissolve  2  c.c.  ammonium  chloride  in  5  c.c.  water,  and 
add  a   piece   of   sodium   amalgam    (Na-J-Hg).     Results? 
The  product  is  ammonium  amalgam.     Note  what  happens 
to  it.     Odor  ?     Reaction  of  solution  ? 

Note.  —  Do  not  throw  away  the  resulting  mercury,  but  ask 
what  to  do  with  it. 

b.  Add  5  c.c.  water  to  3  c.c.  powdered  potassium  chloride 


74  LABORATORY  EXERCISES. 

and  shake  thoroughly.  Pour  off  the  solution  and  add  to  it 
5  c.c.  of  a  concentrated  solution  of  tartaric  acid,  H2C4H4O6. 
Make  this  by  shaking  5  c.c.  powdered  tartaric  acid  with  15 
c.c.  water.  Wait  for  result.  Result  ?  The  product  is  po- 
tassium hydrogen  tartrate.  Equation  ? 

c.  Repeat  #,  using  sodium  chloride  in  place  of  potassium 
chloride.     Result  ?     Repeat  again,  using  ammonium  chlor- 
ide in  place  of  potassium  chloride.     Result  ? 

d.  From  Experiment  XXV,  b  and  c,  tell  what  happens 
when  ammonium  salts  are  treated  with  alkalies.     How  dis- 
tinguish between  sodium  salts  on  the  one  hand  and  ammon- 
ium and  potassium  salts  on  the  other?     Between  sodium 
salts  and  potassium  salts  (two  ways)  ? 

e.  Get  from  the  instructor  two  unknown  substances,  and 
determine  if  they  are  salts  of  sodium,  potassium,  or  ammo- 
nium. 


EXPERIMENT   LVIII. 
CALCIUM. 

Apparatus.  —  Triangle  of  iron  wire,  ring  stand,  blast-lamp, 
evaporating  dish,  platinum  wire,  and  coin. 

Materials.  — Lumps  of  marble,  lime-water,  red  litmus  paper, 
old  mortar,  plaster  of  Paris,  paper,  calcium  chloride,  calcium 
sulphate,  and  ammonium  carbonate  solution. 

a.  Touch  a  piece  of  wet  red  litmus  paper  with  a  piece  of 
marble.  Result?  Support  a  lump  of  marble  about  5  c.c.  in 
volume  on  a  triangle  of  iron  wire  laid  upon  a  ring  of  the 
ring  stand,  and  heat  the  marble  in  the  hottest  Bunsen  flame 
—  in  a  blast-lamp,  if  possible.  When  the  marble  is  cold, 


WATER    OF  CRYSTALLIZATION  IN  GYPSUM.          75 

touch  wet,  red  litmus  with  the  part  that  was  heated. 
Result  ?  Explain.  What  products  are  formed  when  marble 
is  heated  (cf.  Experiment  XL,  b)  ? 

Slake  about  5  c.c.  of  quicklime  by  adding  to  it  water, 
drop  by  drop,  as  long  as  the  water  is  taken  up  readily. 
Wait  for  the  result,  and  describe  it.  Is  there  a  temperature 
effect  ?  Equation  ? 

b.  To  a  piece  of  old  mortar  in  a  test  tube  add  dilute  hy- 
drochloric acid.     Identify  the  gas.     What  does  fresh  mortar 
absorb  from  the  air  ? 

c.  Stir    10    c.c.  plaster  of  Paris  in  an   evaporating   dish 
with   enough   water   to   form   a   fairly   thick   paste.      Put 
the  paste  upon  a  piece  of  paper,  and  push  into  it  a  coin 
slightly  coated  with  vaseline.     At  once  wash  the  evapora- 
ting dish.     Let  the  paste  and  coin  stand  an  hour  or  more. 
Carefully  remove  the  coin  from  the  plaster.     Result  ? 

d.  To  a  solution   containing  a  calcium  salt,  i.  e.,  calcium 
ionS)   add   ammonium   carbonate    solution.     Result?     See 
Experiment  LIV,  d. 

e.  Clean  a  platinum  wire  as  in  Experiment  LIV,  e,  and 
determine  what  color  calcium  chloride  gives  to  the  flame. 
Repeat  with  ^calcium  sulphate.     Be  sure  to  have  concen- 
trated hydrochloric  acid  upon  the  wire. 


EXPERIMENT    LIX. 
WATER  OF  CRYSTALLIZATION  IN   GYPSUM. 

Apparatus.  —  Evaporating  dish,  wire  gauze,  balances, 
Material.  —  Powdered  gypsum  (not  plaster  of  Paris). 


76  LABOEATOEY  EXEECISES. 

a.  Weigh  your  evaporating  dish  (be  sure  it  is  clean  and 
dry),  and  into  it  weigh  accurately  about  3  grams  of  finely 
powdered  gypsum.     Get  the  exact  weight  of  the  gypsum 
taken,  and  record  it. 

b.  Heat  the  evaporating  dish  on  a  clean  wire  gauze  for 
ten  minutes  with  the  hottest  Bunsen  flame.     Then  let  the 
dish   cool,    weigh   it,    and   record    the    result.     Now   heat 
the  dish  again  for  five  minutes,  let  it  cool,  and  determine  the 
weight.     Compare  the  weight  after  the  first   ignition  with 
that  after  the  second.     Keep  your   second  weight  as  the 
final  one. 

c.  Record  your  results  thus  :  — 

Grams. 

Weight  of  evaporating  dish  -|-  gypsum  = 
Weight  of  evaporating  dish  alone  = 
.".  Weight  of  gypsum  taken  = 
Weight  of  evaporating  dish  -f-  calcium  sulphate  = 
Weight  of  evaporating  dish  alone  = 
.'.  Weight  of  water  found  = 
.*.  Per  cent  of  water  in  gypsum  = 


EXPERIMENT   LX. 
STRONTIUM  AND   BARIUM. 

Apparatus.  — Platinum  wire  and  test  tubes. 

Materials.  —  Strontium  chloride  and  nitrate,  barium  chloride 
and  nitrate  ;  solutions  of  strontium  and  barium  chlorides ; 
ammonium  carbonate  solution  ;  dilute  sulphuric  acid. 

a.  Treat  2  c.c.  strontium  chloride  solution  with  a  few 
drops  of  ammonium  carbonate  solution.  Result  ?  Repeat, 


CRYSTAL-WATER  IN  BARIUM  CHLORIDE.          77 

using  barium  chloride  in  place  of  strontium  chloride.    Write 
equations. 

b.  Treat   2  c.c.  strontium    chloride    solution    with    dilute 
sulphuric  acid.    Result?    See  Experiment XXXVI,  a.  Equa- 
tion? 

c.  Clean  the  platinum  wire  as  in  Experiment  LIV,  e,  and 
heat  a  bit  of  strontium  chloride  in  the  flame.     Repeat  with 
strontium  nitrate,  Sr(NO8)2.     Results? 

d.  Repeat  c,  using  the  corresponding  barium  salts.     Re- 
sults ?     How   distinguish   between   calcium,  strontium,  and 
barium  salts  ? 


EXPERIMENT   LXL 
WATER   OF  CRYSTALLIZATION  IN  BARIUM  CHLORIDE. 

Apparatus. — Evaporating    dish,  wire   gauze,  balances,   air 
bath  (?). 

Material. — Barium  chloride,  chemically  pure. 

a.  Have  your  evaporating  dish  clean  and  dryland  get  its 
weight.     Then  weigh  into  it  accurately  about  3  grams  of 
barium    chloride ;    this    should    be    pure,    dry,    and    finely 
powdered. 

b.  Heat  the  evaporating  dish  with  its  contents  in  an  air 
bath  at  120°  to  130°  C.  for  one  hour,  then  cool  it  ten  minutes, 
and  get  its  weight.     Record  your  results  as  in  Experiment 
LIX,  c,  and  get  the  per  cent  of  water  in  the  crystallized 
barium  chloride. 


78  LABOEATOEY  EXEECISEti. 

EXPERIMENT   LXII. 
MAGNESIUM, 

Apparatus.  —  Tongs,  test  tubes. 

Materials.  — Magnesium  wire,  dilute  hydrochloric  acid,  solu- 
tions of  magnesium  sulphate  and  chloride,  disodium  hydrogen 
phosphate,  and  ammonium  chloride  and  Iwdroxide,  magnesite. 

a.  Hold  a  piece  of  magnesium  wire  2   cm.  long  in  the 
flame  (use  tongs).     Result'?     Describe  the  product. 

b.  Treat  a  second  piece  of  magnesium  with  dilute  hydro- 
chloric  acid.      Result?     Identify  the   gas,  and   write   the 
equation.     See  Experiment  VI. 

c.  To  2  c.c.  of  magnesium  sulphate  solution  add  sodium 
carbonate    solution.      Result?      Repeat,    using   magnesium 
chloride  instead  of  the  sulphate. 

d.  See  Experiment  XLVII,  d,  for  the  action  of  a  solution 
containing  a  magnesium  salt  with  disodium  hydrogen  phos- 
phate  and   ammonium   hydroxide.     Rewrite  the   equations 
here. 

Repeat  that  experiment  with  magnesium  chloride  solution 
instead  of  the  sulphate.  Equation  ? 

e.  Treat  a  small   piece   (half   a  c.c.)   of  magnesite  with 
dilute  nitric  acid.     Result  ?     Identify  the  gas,  and  write  the 
equation. 

From  Experiment  XL,  #,  tell  the  effect  of  heat  upon 
magnesite. 


ZINC.  79 


EXPERIMENT   LXIII. 
ZINC. 

Apparatus. —  File  or  sand-paper,  knife,  iron  dish  with  flat 
bottom,  test  tubes,  mouth  blowpipe. 

Materials.  — Zinc,  tin,  lead,  and  copper  ;  zinc  dust ;  solutions 
of  zinc  sulphate,  sodium  hydroxide,  and  ammonium  sulphide  ; 
dilute  sulphuric  and  hydrochloric  acids  ;  hydrogen  sulphide  ; 
stick  of  charcoal. 

a.  Clean  part  of  a  piece  of  zinc  with  a  file  or  with  sand- 
paper.    Color?     Is  zinc  hard  or  soft  (use  a  knife  or  rough 
edge    of  glass)  ?     Place   a  burner  below  the  center  of   an 
iron  dish.     At  equal  distances  from  the  center  place  pieces 
of  zinc,  tin,  lead,  and  copper,  and  determine  the  order  in 
which  they  melt.     Return  the  metals  to  the  proper  bottles. 

b.  Heat  a  piece  of  zinc  on  charcoal  with  the  oxidizing 
flame  produced  by  the  mouth  blowpipe.     Results  ?     To  do 
this  proceed  as  follows :  — 

Hollow  out  a  depression  near  one  end  of  the  charcoal,  and 
into  it  put  the  zinc.  To  make  the  blowpipe  flame,  have  a 
luminous  Bunsen  flame  4  cm.  high,  and  hold  the  blowpipe  so 
that  the  flame  produced  will  be  inclined  about  30  degrees  to 
the  horizontal  plane. 

To  make  an  oxidizing  flame,  hold  the  end  of  the  blowpipe 
inside  the  luminous  flame,  a  centimeter  above  the  tip  of  the 
dark,  inner  cone.  Hold  the  charcoal  at  such  a  distance  that 
the  zinc  is  in  the  outer,  faintly-luminous  part  of  the  blowpipe 
flame. 

To  make  a  reducing  flame,  hold  the  tip  of  the  blowpipe  just 
at  the  outer  edge  of  the  flame  at  its  middle  part,  and  hold  the 
assay  (here,  zinc)  much  nearer  the  blowpipe  than  in  the  oxi- 


80  LABORATORY  EXERCISES. 

dizing  flame.     The  proper  region  is  just  at  the  tip  of  the  inner, 
light-blue  cone  of  the  blowpipe  flame. 

c.  What  action  has  hydrochloric  acid  upon  zinc?     Equa- 
tion ?     See  Experiment  V  for  the  action  of  dilute  sulphuric 
acid,  and  Experiment  X  for  the  behavior  of  zinc  sulphate 
crystals,  ZnSO4.  7  H2O,  when  heated. 

d.  Mix  1  c.c.  zinc  dust  with  5  c.c.  sodium  hydroxide  so- 
lution, and  heat  carefully.     Test  evolved  gas  with  a  flame. 
Result?     The  solution  contains  sodium  zincate,  Na2ZnO2. 
Write  the  equation. 

e.  To  2  c.c.  zinc  sulphate  solution  add  a  drop  of  sodium 
hydroxide  solution.     Result  ?     What,  probably,  is  the  pre- 
cipitate ?      Equation  ?      Repeat   with    a    second    test   tube. 
Now  add  to  the  first  tube  dilute  hydrochloric  acid,  shaking. 
To   the   second  tube   add   an   excess  of  sodium  hydroxide, 
shaking.     Results?     The  alkaline  solution  contains  sodium 
zincate.     Equations  ? 

What  do  these  experiments  show  as  to  the  nature  of  zinc 
hydroxide  ? 

f.  To  10  c.c.  zinc  sulphate  solution  add  a  drop  of  dilute 
sulphuric  acid,  and  then  hydrogen  sulphide.     Result?     Tut 
the  solution  into  a  beaker  and  add  5  c.c.  ammonium  sul- 
phide   solution,    stirring.      Result?      The    product   is    zinc 
sulphide,  ZnS.     Color?     Equation? 

Add  10  c.c.  water,  stir  the  mixture,  let  it  settle,  and  then 
pour  off  the  supernatant  liquid.  Add  15  c.c.  more  water* 
stir,  let  settle,  and  decant ,  i.  e.,  pour  off  the  water.  This  is 
called  "  washing  by  decantation." 

To  the  zinc  sulphide  add  dilute  sulphuric  acid.  Result? 
What  is  the  gas?  Equation?  Why  was  not  zinc  sulphide 
precipitated  by  hydrogen  sulphide? 


CADMIUM.  81 

EXPERIMENT   LXIV. 
EQUIVALENT   OF  ZINC. 

Apparatus.  —  Same  as  in  Experiment  VI. 
Materials.  — Zinc,  in  sheet  form  or  in  sticks  ;  dilute  (5%)  sul- 
phuric acid. 

a.  Dissolve  zinc  in  dilute  sulphuric  acid  just  as  you  did 
magnesium  in  Experiment  VI,  and  find  the  volume  of  hy- 
drogen liberated  by  a  known  weight  of  zinc.     Use  from 
0.45  gram  to  0.55  gram  of  zinc.     If  the  zinc  is  in  sheet 
form,  it  will    react   readily;    but   a   little  impurity,   chiefly 
carbon,  will  remain  insoluble.     If  the  zinc  is  pure,  it  will 
react  with  difficulty ;  therefore  wind  about  the  zinc  a  piece 
of  platinum  wire  or   a  narrow  strip  of  platinum  foil.     Set 
the  apparatus  aside  until  the  zinc  is  in  solution ;  then  pro- 
ceed as  in  Experiment  VI. 

b.  Reduce  the  volume  of   hydrogen  to  standard  condi- 
tions, and  calculate  the  weight  of  the  hydrogen  obtained. 
Finally,  solve  for  x  in  the  proportion,  — 

Weight  of  zinc  taken  :  weight  of  hydrogen  obtained  :: 
x:  1. 

The  value  of  x  will  be  the  equivalent  of  zinc. 


EXPERIMENT   LXV. 
CADMIUM. 

Maierials.  —~  Cadmium    sulphate,   hydrogen    sulphide,    am- 
monium sulphide. 


82  LABORATORY  EXERCISES. 

a.  Dissolve  completely  not  more  than  1  c.c.  cadmium 
sulphate,  CdSO4,  in  5  c.c.  water,  and  add  hydrogen  sulphide 
in  excess.  Result?  The  visible  product  is  cadmium  sul- 
phide^ CdS.  Color?  Equation?  What  other  sulphides 
of  the  same  color  have  you  had  ?  Wash  the  precipitate  by 
decantation,  and  treat  it  with  5  c.c.  ammonium  sulphide. 
Result?  How  distinguish  between  cadmium  sulphide  arid 
other  sulphides  of  the  same  color  ? 


EXPERIMENT   LXYL 
MERCURY. 

Apparatus. — Pipette  (medicine  dropper). 

Materials. — Mercury,  concentrated  nitric  acid,  hydrogen  sul- 
phide, hydrochloric  acid,  sodium  hydroxide  and  potassium 
iodide  solutions,  ammonium  hydroxide,  zinc,  and  copper. 

Caution. — Before  working  with  mercury  remove  all  rings. 
Do  not  throw  mercury  away  ;  but  ask  what  you  are  to  do  with  it. 

a.  By  means  of  a  pipette  take  from  the  mercury  bottle  a 
globule  three  times  as  large  as  an  ordinary  water  drop ;  add 
to  it  2  c.c.  water  and  2  c.c.  concentrated  nitric  acid.     Result  ? 
Let  stand  until  action  stops ;  this  may  take  some  hours. 

b.  While  waiting  for  a,  dissolve  a  globule  of  mercury  the 
size  of  a  water  drop  in  concentrated  nitric  acid ;  this  gives 
-mercuric   nitrate,    Hg(NO3)2.      Equation    (cf.    Experiment 

XXIX,  g)  ?     Dilute  with  15  c.c.  water. 

c.  To  2  c.c.  mercuric  nitrate  solution  (b)  add  hydrogen 
sulphide.     Result?     The  precipitate  is  mercuric  sulphide, 
HgS,    Equation  ? 


MEECUEY. 


83 


d.  Add  to  separate  portions  of  the  nitrate  solution,  hydro- 
chloric acid,   sodium  hydroxide  solution,  and   potassium 
iodide  solution,  respectively.     Results?     Add  the  potassium 
iodide    drop   by   drop,   noting   changes.     Write   equations 
where  possible. 

Note.  —  With  sodium  hydroxide  we  should  expect  mercuric 
hydroxide,  Hg(OH)2  ;  this,  however,  decomposes  into  the  oxide 
and  water. 

e.  Note  the  result  of  a.     The  crystals  are  mercurous  ni- 
trate, HgNO3 ;  pour  out  into  a  beaker,  and  add  15  c.c.  water 
and  a  drop  of  strong  nitric  acid. 

/.  To  2  c.c.  of  the  mercurous  nitrate  solution  of  e  add  hy- 
drogen sulphide.  The  precipitate  is  mercuric  sulphide  and 
mercury.  Write  the  equation. 

g.  Repeat  d  with  the  mercurous  instead  of  the  mercuric 
nitrate.  Results  ?  With  sodium  hydroxide  the  precipitate 
is  mercurous  oxide,  Hg2O.  Write  the  equations.  Treat  the 
precipitate  produced  by  hydrochloric  acid  with  ammonium 
hydroxide.  Result  ? 

h.  Into  the  rest  of  the  mercurous  nitrate  put  a  strip  of 
zinc  and  a  copper  wire.  Results?  Now  rub  them  dry. 
Results  ? 

i.  Classify  the  results  of  c,  d,  f,  and  g  in  five  vertical 
columns. 


Formula  of 
Precipitant. 

Mercuric  Nitrate. 

Mercurous  Nitrate. 

Formula  of 
Ppt. 

Color. 

Formula  of 
Ppt. 

Color. 

NaOH,  etc. 

84  LABORATORY  EXERCISES. 

EXPERIMENT   LXVIL 
COPPER. 

Apparatus.  —  File  or  sand-paper. 

Materials.  —  Copper  wire,  concentrated  hydrochloric  acid; 
solutions  of  cupric  sulphate,  ammonium  hydroxide,  sodium  hy- 
droxide, and  cupric  nitrate  ;  grape-sugar  ;  iron  nail. 

a.  File  a  piece  of  copper  bright.     Color  ?     Is  it  hard  or 
soft  ?     From  Experiment  LXIII  give  its  fusibility  compared 
with  zinc,  etc.     By  holding  one  end  of  the  wire  in  the  flame 
determine  if  it  is  a  conductor  of  heat. 

b.  From  Experiments  XXIX  and  XXXIV  tell  the  action 
of  nitric  and  sulphuric  acids  upon  copper.     Find  out  if  cop- 
per reacts  readily  with  concentrated  hydrochloric  acid. 

c.  To  2  c.c.  cupric  sulphate  solution  add  ammonium  hy- 
droxide solution  in  excess.     Result?     Repeat  with  sodium 
hydroxide  instead  of  ammonium  hydroxide.     Result?     Re- 
peat, having  the  cupric  sulphate  hot,  and  then  add  the  sodium 
hydroxide.     Result?     The  last  precipitate  is  cupric  oxide, 
CuO.     How  formed  (cf.  Experiment  LXVI,  d  and  g)  ? 

d.  From  Experiment  XXXIII,  b  and  c,  tell  the  effect  of 
hydrogen  sulphide  upon  cupric  sulphate.     Equation?     Pass 
hydrogen    sulphide    into    cupric    nitrate    solution.     Result? 
Equation?     What  is  the  effect  of  heating  Hue  vitriol  (cf. 
Experiment  X,  d)  ? 

e.  Dissolve  half  a  c.c.  powdered  grape-sugar,  C6H12O6,  in 
2  c.c.  water,  and  add  it  to  5  c.c.  cupric  sulphate  solution. 
Now   add    sodium    hydroxide   solution,   shaking  until   the 
precipitate    first    formed    is    redissolved,     Color?    Warm 


SILVER.  85 

carefully,  noting  changes.  Let  stand.  Results?  Color 
of  product?  It  is  cuprous  oxide,  Cu2O.  What  effect  had 
the  grape- sugar  ? 

f.  Put  an  iron  nail  into  cupric  sulphate  solution.    Result  ? 

EXPERIMENT   LXVIIL 
SILVER. 

Materials. —  Silver  foil,  silver  nitrate  solution,  nitric  acid,  sod- 
ium thiosulphate  ;  solutions  of  sodium  chloride  and  potassium 
bromide,  iodide,  and  cyanide  ;  filter  paper;  hydrogen  sulphide. 

a.  In  a  test  tube  treat  a  piece  of  silver  foil  with  2  c.c. 
concentrated  nitric  acid.     Result  ?     Equation  ?     Dilute  with 
water  to  10  c.c. 

b.  To  2  c.c.  of  the  solution  of  a  add  5  c.c.  sodium  chloride 
solution.     Result  ?     Equation   (cf.    Experiment  XVII,  j)  ? 
Boil  the  contents  of  the  tube.     Result  ?     Get  the  precipitate 
on  filter  paper,  and  expose  it  to  sunlight.     Result  ? 

c.  To  5  c.c.  of  solution  a  add  1  c.c.  potassium  bromide  so- 
lution.    Result?     Heat  to  boiling,  pour  off  the  supernatant 
liquid,  and  add  to  half  of  the  precipitate  sodium  thiosul- 
phate solution,  Na2S2O3  (make  this  by  dissolving  the  crystals 
in  water).     Result?     Expose  the  other  half  on  filter  paper 
to  sunlight.     Result? 

d.  To  1  c.c.  silver  nitrate  solution  add  1   c.c.  potassium 
iodide  solution.     Result?     Equation? 

e.  To  Ice.  silver  nitrate  solution  add  hydrogen  sulphide. 
Result  ?     Equation  ? 

f.  To  1  c.c.  silver  nitrate  solution  addpotassium  cyanide 


86  LABOEATOEY  EXEECI8ES. 

solution,  drop  by  drop.  Result?  Equation?  Continue 
adding  it,  shaking,  until  it  is  in  excess.  Result?  The  so- 
lution contains  the  double  cyanide,  KCN.AgCN,  i.  e.y 
KAg(CN)2.  Add  sodium  chloride  solution.  Result?  Ex- 
plain the  result  in  terms  of  the  ionic  theory  (cf.  Experi- 
ment LII). 


EXPERIMENT   LXIX. 
ALUMINUM. 

Apparatus.  —  Test  tubes,  tongs,  blast-lamp. 

Materials. — Aluminum  wire  and  filings,  white  muslin,  hy- 
drochloric acid  ;  solutions  of  sodium  hydroxide,  aluminum  sul- 
phate, sodium  carbonate,  alum,  ammonium  hydroxide,  and 
cochineal ;  powdered  alum,  sodium  bicarbonate,  potassium  sul- 
phate, ammonium  sulphate,  aluminum  sulphate. 

a.  Determine  whether  aluminum  is  a  conductor  of  heat 
as  in  Experiment  LXVII,  a.     Does  the  wire  melt  in  the  Bun- 
sen  flame  (use  tongs)  ?     Try  the  blast-lamp.     Result  ? 

b.  To  2  c.c.  aluminum  filings  add  5  c.c.  concentrated  hy- 
drochloric acid,  and  warm.     Result  ?     Test  the  gas.     Equa- 
tion? 

c.  Wash  the  filings  remaining  from  #,  by  decantation,  add 
5   c.c.   concentrated  sodium  hydroxide  solution,  and  warm 
carefully.     Determine  the  nature  of  the  gas  evolved.     Re- 
sult?    The  solution  contains  sodium  aluminate,  Na3AlO3 
(cf.  Experiment  LXIII,  d).     Equation? 

d.  To  5  c.c.   of   aluminum   sulphate  solution,  A12(SO4)3, 
add  1  c.c.  sodium  hydroxide  solution.     Result?     Equation? 
Get  half  of  the  precipitate  into  a  second  test  tube,  and  add 


IKON.  87 

an  excess  of  sodium  hydroxide  solution.  Result?  If  the 
solution  now  contains  sodium  aluminate,  Na3AlO3,  write 
the  equation.  To  the  other  half  of  the  precipitate  add  hy- 
drochloric acid.  Result?  Equation?  Compare  with  this 
the  behavior  of  zinc  hydroxide. 

e.  Dissolve  as  much  ammonium  sulphate  as  possible  in 
5  c.c  hot  water,  and  add  to  it  in  a  beaker  5  c.c.  water  simi- 
larly saturated  with  aluminum  sulphate.     Cool  the  mixture. 
Result  ?     The  product  is  ammonium  alum.     Heat  again  to 
complete  solution,  and  let  stand  over  night.    Result  ?   Shape 
of  crystals  ? 

f.  Repeat  e,  using  potassium  sulphate   instead   of   am- 
monium sulphate.     Results  ?     Compare  the  crystals. 

g.  To  5  c.c.  of   the  solution  of  any  aluminum  salt  add 
sodium   carbonate  solution.     Result  ?     Identify  the  escap- 
ing gas.     The  precipitate  is  aluminum  hydroxide,  A1(OH)3. 
Mix  a  cubic  centimeter  of  powdered  alum  with  a  cubic  centi- 
meter of  sodium  bicarbonate,  and  add  water.     Result  ? 

h.  To  1  c.c.  of  a  solution  of  cochineal  add  5  c.c.  alum  so- 
lution, immerse  a  piece  of  white  muslin,  and  then  add  am- 
monium hydroxide  solution,  shaking.  Results  ? 


EXPERIMENT   LXX. 
IRON. 

Apparatus.  —  Test  tubes,  tongs,  blast-lamp,  magnet,  beaker. 

Materials.  — Iron  wire  and  filings,  copper  wire  ;  hydrochloric, 

sulphuric,  and   nitric   acids ;   solutions  of  potassium   ferrocy- 


88  LABORATORY  EXERCISES. 

anide,  ferricyanide,  and  sulphocyanate  ;  ammonia  water,  hy- 
drogen sulphide,  ammonium  sulphide,  solid  ferrous  sulphate, 
and  ferric  chloride. 

a.  Compare  the  heat  conductivity  of  iron  wire  with  that 
of  copper.    Test  its  magnetic  properties  ;  its  fusibility  in  the 
Bunsen  flame  and  the  blast- lamp.     Results  ? 

b.  Treat  3  c.c.  iron  filings  in  a  beaker  with  20  c.c.  dilute 
hydrochloric  acid,  stirring.     Results  ?     Identify  the  gas.     If 
the  solution  contains  ferrous  chloride,  Fe012,  write  the  equa- 
tion.    When  action  almost  ceases,  filter  off  10  c.c.  of  the 
solution.     Color  of  filtrate? 

c.  Divide  the  filtrate  of  b  into  four  parts.     To  the  first 
add    a    few    drops    of    potassium   ferricyanide    solution, 
K,Fe(CN)6.     Result?    This  is  "Turn&ulFs  Hue?     To  the 
second  portion    add    ammonia- water.     Result?     Equation? 
Note  any  change  on  standing  in  the  air.     To  the  third  part 
add  potassium  f err  ocy  anide,  K4Fe(CN)6.     Result?    To  the 
last  portion  add  potassium  sulphocyanate  solution,  KSCN", 
Result  ? 

Wash  out  your  test  tubes  and  beakers  at  once. 

d.  Filter  the  remainder  of  the  ferrous  chloride  solution  ol 
#,  and  add  2  c.c.  concentrated  nitric  acid.    Heat  carefully  for 
two  minutes  in  a  beaker.     Resulting  color  ?     The  solution 
contains  ferric  chloride  and  nitrate.     To  a  drop  of  it  in  a 
test  tube  add  a  drop  of  potassium  ferricyanide  solution ;  if 
it  still  gives  a  blue  precipitate,  add  2  c.c.  more  nitric  acid, 
and  boil  again. 

Treat  the  resulting  substance  in  four  test  tubes  with  the 
reagents  used  in  c.  Result  in  each  case  ? 

The  precipitate  from  potassium  ferrocyanide  and  a  ferric 
salt  is  "Prussian  blue." 


IE  ON. 


e.  Classify  the  results  of  c  and  d  (last  part)  in  five  ver- 
tical columns. 


Formula 
of  Keagent. 

Ferrous  Chloride. 

Ferric  Chloride. 

Precipitate 
or  Solution? 

Color. 

Precipitate 
or  Solution? 

Color. 

f.  In  a  test  tube  shake  2  c.c.  powdered  ferrous  sulphate 
with  10  c.c.  water,  pour  off  half  of  the  solution,  and  pass 
hydrogen  sulphide  into  it.  Result?  Does  all  the  iron 
appear  to  be  precipitated?  Write  the  equation  represent- 
ing the  reaction  you  would  expect  to  take  place. 

From  Experiment  XXXIII  tell  the  effect  of  dilute  sul- 
phuric acid  upon  ferrous  sulphide.  Write  the  equation 
here.  Compare  these  two  equations.  Conclusion? 

To  the  other  half  of  the  ferrous  sulphate  solution  add  five 
drops  of  dilute  sulphuric  acid,  and  pass  in  hydrogen  sulphide. 
Compare  with  the  result  without  the  acid  ?  Now  add  am- 
monium sulphide.  Result  ?  Equation  ? 

g.  Dissolve  1  c.c.  ferric  chloride,  FeCl3,  in  10  c.c.  water, 
and  pass  in  hydrogen  sulphide  at  least  two  minutes.  Re- 
sult? Boil  the  contents  of  the  tube,  and  then  filter.  Test 
the  filtrate  with  a  drop  of  potassium  ferricyanide  solution. 
Result  and  conclusion  ? 

Determine  the  nature  of  the  residue  on  the  filter  paper  by 
collecting  it  on  a  piece  of  porcelain  and  igniting  it.  Odor  ? 

Write  the  equation  for  the  action  of  hydrogen  sulphide 
on  ferric  chloride. 


90  LABORATORY  EXERCISES. 

EXPERIMENT  LXXL 
NICKEL  AND  COBALT. 

Apparatus.  — Platinum  wire,  test  tubes. 

Materials.  —  Nickel  and  cobalt  and  their  nitrates  ;  solutions 
of  the  nitrates  ;  borax,  sodium  hydroxide  solution,  concen- 
trated, chemically  pure  hydrochloric  acid. 

a.  Give  the  physical  properties  of  cobalt  and  nickel  from 
an  examination  of  the  metals.     Effect  of  a  magnet  ? 

b.  To  2  c.c.  nickel  nitrate  solution,  Ni(NO3)2,  add  a  drop 
of  hydrochloric  acid  and  then  hydrogen  sulphide.     Result  ? 
Now  add  ammonium  sulphide.     Result  ?     Equation  ?     Ex- 
plain the  results  from  Experiments  XXXIII  and  LXX,  f. 

c.  Make  a  borax  bead  as  in  Experiment  LI,  a  and  ft,  and 
determine  the  color  given  to  it  by  nickel  nitrate. 

d.  Repeat  b  and  c  with  cobalt  nitrate,  Co(NO3)2,  instead 
of  nickel  nitrate.     Results  ? 

e.  To  2  c.c.  cobalt  nitrate  solution  add  sodium  hydroxide 
solution,  a  drop  at  a  time,  until  it  is  in  excess.     Results  ? 

f.  To  2  c.c.  cobalt  nitrate  solution  add  5  c.c.  concentrated, 
chemically  pure  hydrochloric   acid.     Result?     Dilute  with 
water.     Result  ? 


EXPERIMENT   LXXII. 
MANGANESE    COMPOUNDS. 

Apparatus. — Platinum  wire,  test  tubes. 

Materials. — Manganese  sulphate,  potassium  permanganate, 


CHROMIUM  COMPOUNDS.  91 

ferrous   sulphate,  grape-sugar,  ammonia-water,  hydrogen   sul- 
phide, and  ammonium  sulphide. 

a.  Dissolve  1  c.c.  powdered  manganese  sulphate,  MnSO4, 
in  5  c.c.  water.     To  half  of   it  add  a  drop  of  dilute  sul- 
phuric acid  and  then  hydrogen  sulphide.      Result?      Now 
add  ammonium    sulphide.      Result  ?      Color  ?      Equation  ? 
Explain  the  results. 

b.  To  the  other  half  of  solution  a  add  ammonia-water. 
Result  ?     Eq  nation  ? 

c.  To    2  c.c.    ferrous    sulphate   solution  (cf.  Experiment 
LXX,  /)   add  potassium  permanganate  solution.     Result? 
Continue,  drop  by  drop,  until  the  solution  is  just  faintly 
pink.     Now  add  ammonia-water.     State   and   explain  the 
result  (cf.  Experiment  XXXIV,  /). 

d.  Dissolve  a  crystal  of  potassium  permanganate,  KMnO4, 
in  wrater,  and  add  grape-sugar  solution.     Result?     Explain 
(cf.  Experiment  LXVII,  e). 

e.  From  Experiment  XVI  tell  the  action  of  manganese 
dioxide  with  hydrochloric  acid ;  from  Experiment  VII,  with 
potassium  chlorate ;  from  Experiment  XL VI,  with  hydro- 
gen peroxide /  and  from  Experiment  LI,  a,  tell  the  color  of 
the  manganese  bead. 


EXPERIMENT   LXXIII. 
CHROMIUM   COMPOUNDS. 

Apparatus.  — Chlorine  generator,  platinum  wire,  evaporating 
dish,  test  tubes. 

Materials.  — •  Solutions  of  potassium  chromate  and  dichrom- 


92  LABOR  A  TOE  Y  EXERCISES. 

ate,  of  chromic  chloride,  of  potassium  and  sodium  hydroxides; 
hydrochloric  acid,  alcohol,  borax,  barium  chloride  solution, 
chrome-alum. 

a.  What   is    the    color    of    solutions   of   potassium   di- 
chromate  (K2Cr2O7),  potassium  chromate  (K2CrO4),  and  of 
chromic  chloride  (Cr013)  ? 

b.  Treat  1  c.c.  of  potassium  dichromate  solution  with  a 
drop  of  potassium  hydroxide  solution.     Result?     From  the 
color  tell  what  is  formed. 

Complete  the  equation, 

K2Cr207  +  2  KOH »  ?  +  ? 

c.  To  1  c.c.  potassium  chromate  solution  add  a  drop  of 
concentrated  hydrochloric  acid.     Result  ?     What  is  formed  ? 

2  K2CrO4  +  2  HC1 »  ?  +  ?  -f  ? 

How  can  a  dichromate  be  changed  to  a  chromate  ?     A 
chromate  to  a  dichromate  ? 

d.  To    2  c.c.   potassium    chromate    solution   add    barium 
chloride    solution.       Result?       Equation?       Repeat    with 
chromic   chloride   instead    of    the    chromate.       Result? 

e.  To    1   c.c.  chromic    chloride    solution    add   a    drop  of 
sodium    hydroxide    solution.      Result?      Equation?      Now 
add  the  alkali  in  excess,   shaking.     Result?     The  solution 
contains  a  chromite,  NaCrO2.     What  other  elements  behave 
in  this  way  ?     See  Experiments  LXIII  and  LXIX.     Save 
for  g. 

f.  Repeat  e  with  potassium  chromate  instead  of  chromic 
chloride.     Result  ?     In  what  three  ways  can  a  chromic  salt 
be  distinguished  from  a  chromate? 

g.  To  the  clear  solution  of  e  add  chlorine  gas  until  there 


LEAD.  93 

is  no  further  change.  Do  this  in  a  gas-chamber.  Results  ? 
Test  the  resulting  solution  as  in  b,  c,  and  d.  Results  ?  How 
can  a  chromic  salt  be  changed  into  a  chromate  ? 

A.  To  10  c.c.  potassium  dichromate  solution  in  an  evapo- 
rating dish  add  2  c.c.  concentrated  hydrochloric  acid  and 
2  c.c.  ethyl  alcohol.  Boil  until  bright  green,  but  not  to  dry- 
ness.  Test  a  part  of  the  liquid  with  barium  chloride  solu- 
tion. Result  ?  With  potassium  hydroxide  solution.  What 
does  the  green  solution  contain?  How  can  a  chromate 
be  changed  to  a  chromic  salt?  See,  also,  Experiment 
XXXIV,  / 

i.  Refer  to  Experiment  LI  for  the  borax  bead  test.  Re- 
peat with  a  tiny  piece  of  chrome-alum.  Result  ? 


EXPERIMENT    LXXXIV. 
LEAD. 

Apparatus.  — File  or  knife,  test  tubes,  mouth  blowpipe. 

Materials.  —  Lead  ;  hydrochloric,  nitric,  and  sulphuric  acids  ; 
lead  nitrate,  solutions  of  potassium  chromate  and  sodium  hy- 
droxide, lead  oxide,  stick  of  charcoal. 

a.  File  or  cut  off  the  coating  on  lead.     Is  it  hard  or  soft  ? 
Color?     Try  to  mark  on  paper  with  lead.     Result?     Refer 
to  Experiment  LXIII,  a,  for  its  fusibility.     Treat  a  small 
piece  with  hydrochloric  acid,  both  the  dilute  and  the  strong. 
Results  ?     Wash  the  lead,  and  add  2  c.c.  concentrated  nitric 
acid  and  2  c.c.  water.     Heat  gently.     Result?     Write  the 
equation  (cf.  Experiment  XXIX). 

b.  Heat  one-fourth  of  a  c.c.  of  lead  monoxide  on  charcoal 


94  L  AS  OH AT OE  Y   EXERCISES. 

in  the  reducing  flame  (mouth  blowpipe).     See  Experiment 
LXIII,  b.     Result?     How  identify  the  product  ? 

c.  Dissolve   2  c.c.  powdered  lead  nitrate,  Pb(NO3)2,  in 
15  c.c.  water,  heating.     Cool,  and  add  to  2  c.c.  of  the  solu- 
tion 5  c.c.   dilute   hydrochloric    acid.     Result?     Equation? 
Wash  the  precipitate  by  decantation,  and  heat  it  with  10 
c.c.  water.     Result  ?     Cool  the  solution.     Result  ? 

d.  To  2  c.c.  of  the  lead  nitrate  solution  add  dilute  sul- 
phuric acid.     Result?     Use  potassium   chr ornate   solution 
instead  of  sulphuric  acid.     Result  ?     Equation  in  each  case  ? 

From  Experiment  XXXIII,  d,  tell  effect  of  hydrogen 
sulphide  upon  lead  nitrate.  For  the  reduction  of  lead  oxide 
by  charcoal,  see  Experiment  XL,  a. 

e.  To  2  c.c.  lead  nitrate  solution  add  a  drop  of  sodium 
hydroxide   solution.     Result?      Equation?      Now   add   an 
excess,    shaking.     Result?     What   three   other   hydroxides 
behaved  in  the  same  way  ?     See  Experiment  LXXIII,  e. 

f.  Put  into  the  remainder  of  the  lead  nitrate  solution  a 
strip    of    zinc.     Leave   it   at   least   ten   minutes.     Result? 
Equation  (cf.  Experiment  LXVII,  /)  ? 


EXPERIMENT  LXXV. 
TIN. 

Apparatus. — Test  tubes,  stopper  and  deliver}^  tube,  mouth 
blowpipe. 

Materials. — Tin  (granular  and  in  a  bar)  ;  concentrated  hy- 
drochloric and  nitric  acids  ;  solutions  of  mercuric  chloride, 
stannic  chloride,  and  sodium  hydroxide  ;  ammonium  sulphide, 
hydrogen  sulphide,  zinc,  sulphur,  stick  of  charcoal. 


TIN.  95 

a.  Treat  about  2  c.c.  of  small  bits  of  tin  with  10  c.c.  con- 
centrated hydrochloric  acid  in  a  test  tube.    Warm  gently  to 
start   the  action,  and   when   the  effervescence  is  vigorous 
attach  a  stopper  and  delivery  tube  and  collect  the  gas  over 
water.     Identify  the  gas.     Result?     The  solution  contains 
stannous  chloride,  SnCl2.     Equation  ?     Let  the  action  con- 
tinue at  least  ten  minutes. 

b.  From  Experiment  LXIII,  a,  compare  the  fusibility  of 
tin  with  that  of  lead,  etc.     Hold  a  bar  of  tin  near  your  ear, 
and  bend  it.     Result  ?     What  color  has  bright  tin  ?     Is  it 
hard  or  soft  ? 

c.  To  1  c.c.  mercuric  chloride  solution,  HgCl2,  add  4  or  5 
c.c.  of  your  stannous  chloride  solution,  and  then  heat.     Note 
all  the   changes.     The   solution  contains    stannic   chloride, 
SnCl4.    Equation  ? 

d.  To  2  c.c.  stannous  chloride  solution  add  5  c.c.  water 
and  then  hydrogen  sulphide.     Result?     Color?     Equation? 
Wash  the  precipitate   by  decantation,  and  add  5  c.c.   am- 
monium sulphide  (use  an  evaporating  dish  or  beaker)  and  a 
small  lump    of   sulphur.     Warm  gently,  and  stir.     Result? 
Cool,  and  add  dilute  hydrochloric  acid  in  excess.     Result  ? 
Compare  the  color  with  that  of  the  original  precipitate. 

e.  To  2  cX5.  stannous  chloride  solution  add  1  c.c.  concen- 
trated nitric  acid,  and  heat  gently.     The  solution  contains 
stannic  chloride.    Dilute  with  5  c.c.  water,  and  pass  in  hydro- 
gen sulphide.     Result?     Color?     Stannic  sulphide  is  SnS2. 
Equation?     Wash  the  precipitate  by  decantation,  add  am- 
monium sulphide  and  a  bit  of  sulphur,  and  warm  gently. 
Result?     Add  an  excess  of  dilute  hydrochloric  acid.     Re- 
sult?    Compare  with  the  color  of  the  original  precipitate, 
and  with  that  obtained  at  the  end  of  d.     Conclusion  ? 


96  LABORATORY  EXERCISES. 

f.  To  2  c.c.  stannic  chloride  solution  add  sodium  hydrox- 
ide solution,  drop  by  drop.     Result?     Add  an  excess.     Re- 
sult?    What  other  hydroxides  have  behaved  in  the   same 
way?     See  Experiment  LXXIV,  e. 

g.  Pour  the  solution  of  a  from  any  unused  tin,  and  put 
into  it  a  strip  of  zinc.     Result  ?     Equation  ?     Compare  with 
Experiment  LXXIV,/ 

h.  Heat  a  piece  of  tin  on  charcoal  in  the  oxidizing  flame 
(mouth  blowpipe).     See  Experiment  LXIII,  b.     Results? 


APPENDIX. 


TABLE  OF  EQUIVALENTS  IN  ENGLISH  AND  METRIC 

UNITS. 


A.     LENGTH. 
1  centimeter  — 0.3937  in. 
1  decimeter  —10  cm. 
1  meter          =  100  cm.  — 1,000  mm. 
1  meter          =  39.37  in.  =  3.28  ft. 
1  kilometer   =1,000  m.  =  0.6214  mile. 


1  inch  =2.54  cm. 

1  foot  =  0.3048m. 

1  mile  =1.6094  km. 

B.     AREA. 

1  sq.  cm.       =  0.155  sq.  in. 

1  sq.  m.          =  10.764  sq.  ft.  =1.196  sq.  yd. 

1  sq.  km.       =0.385  sq.  mile. 

C.  VOLUME. 

1  cu.  cm.       =  0.061  cu.  in. 

1  cu.  m/'      =  35.315  cu.  ft. 

1  liter  =  1,000  cu.  cm.  =  1. 0567  qt.  (U.  S.) 

D.  WEIGHT. 
1  gram           =  15.4324  grains. 

1  kilogram     =1,000  grams  =2. 2046  Ibs. 
1  grain  =  0.0648  gram. 

1  ounce  (avoirdupois)  =  28. 35  grams. 
1  ounce  (troy)  =31.1  grams. 

i 


11 


APPENDIX. 


TABLE  OF  ATQMIC  MASSES  AND  SPECIFIC  HEATS.. 


ATOMIC  MASSES.* 


SPECIFIC 


Clarke. 

Richai'ds. 

!  1  I  ,   \  A  O« 

H  =  l. 

1  O  =  16. 

O=16. 

Aluminum  .  . 

.  .  .  Al    ... 

26.9 

27.1 

27.1 

0.214 

Antimony    .  . 

.  .  .  Sb    ... 

319.5 

120.4 

120.0 

0.0508 

Ar^on  

.  .  .  A  .  .  .  . 

39.0 

39.96 

39.92 

Arsenic 

As   . 

74.45 

75.0 

75.0 

0.0814 

Barium  .... 

.  .  .  Ba   ... 

136.4 

137.40 

137.43 

Bismuth    .  .  . 

.  .  .  Bi    ... 

206.5 

208.1 

208.0 

0.0308 

Boron     .... 

.     .  B  .  .  .  . 

10.9 

11.0 

11.0 

0.3G6 

Bromine    .  .  . 

.  .  .  Br   ... 

79.35 

79.95 

79.955 

0.0843-f 

Cadmium  .  .  . 

...  Cd  ... 

111.55 

112.4 

112.3 

0.0567 

Caesium     .  .  . 

.  .  .  Cs    ... 

131.9 

132.9 

132.9 

Calcium  .... 

.  .  .  Ca    ... 

39.8 

40.1 

40.1 

0.170 

Carbon  .... 

.  .  .  C  .  .  .  . 

11.9 

12.0 

12.001 

0.459ft 

Cerium  .... 

.  .  .  Ce    ... 

138.0 

139.0 

140. 

0.0448 

Chlorine    .  .  . 

.  .  .  Cl    ... 

35.18 

35.45 

35.455 

Chromium  .  . 

.  .  .  Cr    ... 

51.7 

52.1 

52.14 

0.100 

Cobalt  

...  Co   ... 

58.55 

59.00 

59.00 

0.107 

Columbiuin    . 

.  .  .  Cb    ... 

93.0 

93.7 

94. 

Copper  .... 

...  Cu  ... 

63.1 

63.60  , 

63.60 

0.0952 

Erbium  .... 

.  .  .  Er    ... 

164.7 

166.0 

166. 

Fluorine    .  .  . 

.  .  .  Fl    ... 

18.9 

19.05 

19.05 

Gadolinium.  . 

...  Gd  ... 

155.2 

156.4 

156.  ? 

Gallium  .... 

...  Ga  ... 

69.5 

70.0 

70.0 

0.079 

Germanium    . 

.  .  .  Ge   ... 

71.9 

72.5 

72.5 

i 

Glucinum    .  . 

.  .  .  Gl    ... 

9.0 

9.1 

9.1 

0.058 

Gold    

.  .  .  Au  .  .  . 

195.7 

197.2 

197.3 

0.0324 

Helium  .... 

.  .  .  He  .  .  . 

3.93 

3.96 

3.96 

Hydrogen    .  . 

.  .  .  H.  .  .  . 

1.000 

1.008 

1.0075 

.  .  .  In    ... 

113.1 

114.0 

114. 

0.0570 

Iodine       . 

.     I         .  . 

125.89 

126.85 

126.85 

0.0541 

Iridium  .... 

.  .  .  Ir.  .  .  . 

191.7 

193.1 

193.0 

0.0326 

Iron 

.  Fe    . 

55.5 

55.9 

55.9 

0.114 

Krypton    .  .  . 

.  .  .  Kr   ... 

81.15 

81.76 

81.7 

Lanthanum    . 

.  .  .  La    ... 

137.6 

138.6 

138.5 

0.0449 

Lead  

.  .  .  Pb   ... 

205.36 

206.92 

206.92 

0.0314 

Lithium    .  .  . 

.  .  .  Li    ... 

6.97 

7.03 

7.03 

0.941 

Magnesium  .  . 

...  Mg  ... 

24.1 

24.3 

24.36 

0.250 

Manganese  .  . 

...  Mn  ... 

54.6 

55.0 

55.02 

0.122 

Mercury    .  .  . 

...  Hg  ... 

198.50 

200.0 

200.0 

0.03  I9f 

*Table  of  Atomic  Masses,  prepared  by  Prof.  F.  W.  Clarke;  "  Journal  of  the 
American  Chemical  Society,"  Vol.  XXIV,  No.  3;  March,  1902. 
t  Solid,    ft  Diamond. 


APPENDIX. 


Ill 


TABLE  OF  ATOMIC  MASSES  AND  SPECIFIC  HEATS.-  Continued. 


ATOMIC  MASSES. 

SPECIFIC 
HEATS. 

0.0722 

0.108 
0.0311 

0.0593 
0.189* 
0.0324 
0.166 

0.0580 
0.0611 

0.0762f 
0.203f 
0.0570 
0.293 

• 

o.nstt 

0.0474 

0.0335 
0.0276 

0.0562 
0.1485 
0.0334 
0.0277 

0.0955 
0.0662 

Molybdenum  .  . 
Neodymium   .  . 
Neon    
Nickel  
Nitrogen 

.  .  Mo  .  .  . 
.  .  Nd  .  .  . 

.  .  Ne   .  .  . 
.  .  Ni   ... 

N 

Clarke. 
H  =  l.           O  =  16. 
95.3               96.0 
142.5             143.6 
19.8               19.94 
58.25             58.70 
13.93             14.04 
189.6             191.0 
15.88             16.000 
106.2             107.0 
30.75             31.0 
193.4             194.9 
38.82             39.11 
139.4             140.5 
102.2             103.0 
84.75             85.4 
100.9             101.7 
149.2  ?          150.3  ? 
43.8               44.1 
78.6               79.2 
28.2               28.4 
107.11            107.92 
22.88             23.05 
86.95             87.60 
31.83             32.07 
181.5             182.8 
126.5             127.7 
158.8             160. 
202.61            204.15 
230.8  ?           232.6  ? 
169.4             170.7 
118.1             119.0 
47.8               48.15 
182.6             184. 
237.8             239.6 
51.0               51.4 
127.               128.0 
171.9             173.2 
88.3               89.0 
64.9               65.4 
89.7               90.4 

Richards. 
0  =  16. 
96.0 
143.6 
19.94 
58.70 
14.04 
190.8 
16.00 
106.5 
31.0 
195.2 
39.14 
140.5 
103.0 
85.44 
101.7 
150.0 
44. 
79.2 
28.4 
107.93 
23.05 
87.68 
32.065 
183. 
127.5  ? 
160. 
204.15 
233. 
171.  ? 
119.0 
48.17 
184. 
238.5 
51.4 
128. 
173. 
89.0 
65.40 
90.6 

Osmium  
Oxvaren  . 

.  .  Os    ... 
.  .  0  .  .  .  . 

Palladium    .  .  . 
Phosphorus    .  . 

.  .  Pd  ... 
.  .  P  .  .  .  . 

Platinum  .... 
Potassium 

.  .  Pt    ... 
K  . 

Praseodymium  . 
Rhodium  .... 
Rubidium  .... 
Ruthenium  .  .  . 
Samarium    .  .  . 
Scandium  .... 
Selenium  .... 
Silicon    
Silver  
Sodium  
Strontium    .  .  . 
Sulphur  

.  .  Pr    ... 
.  .  Rh  .  .  . 
.  .  Rb  .  .  . 
.  .  Ru  ... 

..  Sin  ... 
.  .  Sc    ... 
.  .  Se    ... 
..  Si  .... 
.  .  Ag  .  .  . 
.  .  Na  .  .  . 
.  .  Sr    ... 
.  .  S  .  .  .  . 

Tantalum  .... 
Tellurium    .  .  . 
Terbium    .... 
Thallium  .... 
Thorium   .... 
Thulium    .... 
Tin    .  . 

.  .  Ta   ... 
.  .  Te   .  .  . 
.  .  Tr    ... 
.  .  Tl     ... 
.  .  Th   .  .  . 
.  .  Tin  .  .  . 
Sn 

Titanium  .... 
Tungsten  .... 
Uranium  .... 
Vanadium    .  .  . 
Xenon    
Ytterbium   .  .  . 
Yttrium    .... 

.  .  Ti  .  .  .  . 
.  .  W.  .  .  . 
.  .  Ur  .  .  . 
.  .  V  .  .  .  . 
.  .  Xe  .  .  . 
.  .  Yb  .  .  . 
.  .  Y  .  .  .  . 

Zu 

Zirconium    .  .  . 

.  .  Zr    ... 

*  Yellow. 


t  Crystalline. 


ft  Rhombic. 


IV 


APPENDIX. 


TENSION  OF  AQUEOUS  VAPOR  IN  MM.  OF  MERCURY  (REGNAULT). 


TEMP. 

TENSION. 

TEMP.          TENSION. 

TEMP. 

TKNSION. 

0°C. 

4.6 

11°  C.            9.8 

21°  C. 

18.5 

1 

4.9 

12                 10.4 

22 

19.7 

2 

5.8 

13                 11.1 

23 

20.9 

3 

5.7 

14                 11.9 

24 

22.2 

4 

G.I 

15                 12.7 

25 

23.  G 

5 

G.5 

16                  13.5 

2G 

25.0 

6 

7.0 

17                 14.4 

27 

2G.5 

7 

7.5 

18 

15.4 

28 

28.1 

8 

8.0 

19                  1G.3 

29 

29.8 

9 

8.5 

20 

17.4 

30 

31.6 

10 

9.1 

... 

TABLE  OF  SPECIFIC  GRAVITIES  (WATER  =  1). 


Acetic  ucicl  *      .... 

1.053 

Lead   

11.35 

Alcohol  (ethvl)*    .    .    . 

0.794 

Limestone  

3.2 

2.67 

Lithium  

0  59 

6  72 

1  74 

5.7 

Manganese  .    .    . 

7  2  to  8 

9.8 

Mercury  \   •    •    « 

13.596 

2.63 

Nickel     .... 

S  57 

8.3 

Nitric  acid  (cone  )* 

1.42 

Carbon  (gas)     .... 
Carbon  disulphide  *  . 

1.8 
1.27 

Phosphorus     .... 
Platinum     .... 

1.83 
21  5 

Chloroform  *    .... 

1  5 

Potassium   .              .     . 

0  865 

Coal  (anthracite)  .    .     . 

1.2G  to  1.8 

Silicon     .... 

2  49 

Cobalt  

8.8 

Silver      .... 

10  57 

8.9 

Sodium   .     .    . 

0  97 

Diamond  

3.53 

Sulphur  

2  03 

Ether*     

0.72 

Sulphuric  acid     .     . 

1  84 

(ihos     

2  6  to  3.6 

Tin     

7  29 

Gold     

193 

Water  at  0°  C.      ... 

0999 

Hydrochloric  acid  (cone.)* 

1.22 
0918 

„      „  4.07°  C.     .    . 
,      ,,  100°  C.  . 

1.000 
0958 

Iodine  

495 

(sea)      .     .         . 

1  026 

Iron      ...         ... 

7  8 

Zinc    

6  9  to  7.5 

*  At  15°  C. 


t  At  0°C. 


WEIGHT  (IN  GRAMS)  OF  A  LITER  OF  THE  DRY  GAS  AT  0°  C. 
AND  760  MM. 


Air       

1.293 

0.717 

Ammonia          .         .    . 

0.762 

Nitric  oxide    .... 

1.34 

Carbon  dioxide     ... 

1  977 

1.256 

Carbon  monoxide      .    . 
Chlorine  

1.251 
3  18 

Nitrous  oxide  .... 

1.97 
1.429 

Hydrochloric  acid      .    . 

1.61 
0  0896 

Sulphur  dioxide  .    .    . 

2.87 
0.806 

Hydrogen  sulphide    .    . 

1.542 

APPENDIX. 


TABLES. 

Table  I.  —  HEAT  OF   FORMATION  AND  HEAT  or  SOLUTION 
OF  SOME  SUBSTANCES  IN  KILOGRAM-CENTIGRADE  UNITS. 


NAME. 

Formula. 

Heat  of 
Formation. 

Heat  of 
Solution  in 
Water. 

Ozone     

Oo 

30 

Water  (liquid)     

H2O 

68  4 

Hydrogen  peroxide      .... 
Hydrogen  chloride  

H202 
HC1 

45.2 

22 

17  3 

Hydrogen  bromide  

HBr 

12 

19  9 

Hydrogen  iodide     

HI 

6  1 

19  2 

Hydrogen  sulphide      .... 

H2S 

SO2 

2.7 
71 

4.6 

77 

H2SO4 

193  1 

178 

NH3 

12 

8.4 

Nitrogen  tetroxide 

N,O 

0  Q 

Nitrogen  dioxide 

N024 

7  7 

Nitric  oxide    . 

NO 

21  6 

Carbon  dioxide                            . 

•C&-- 

19  6 

Carbon  disulphide        .... 

^-CO,- 

97 

Carbon  monoxide 

CO 

29 

Phosphorus,  red,  fromyellow  form 
Potassium  hydroxide  .... 
Potassium  carbonate   .... 
Potassium  nitrate    

KOH 
K2C03 
KNO3 

27.3 
103.2 

281. 
119 

13.3 

6.5 

—85 

Sodium  chloride      

NaCl 

97  6 

—  1.2 

Sodium  hydroxide 

NaOH 

102 

9  9 

Sodium  carbonate 

NaoCOo 

272  6 

5  6 

Ammonium  chloride    .... 
Calcium  hydroxide      .... 
Magnesium  hydroxide     .     .     . 
Aluminum  hydroxide  .... 
Ferric  hydroxide     

NH4C1 
Ca  (OH)2 
Mg(OH)2 
A1(OH)3 
Fe  (OH)3 

75.8 
215. 
217. 
297. 
198 

—4. 
3. 

Ferrous-ferric  oxide     .... 
Zinc  oxide  

Fe304 
ZnO 

265. 
86 

Cupric  oxide  ....... 

CuO 

37 

Mercuric  oxide   

II  gO 

20.7 

Silver  oxide    

AffoO 

6. 

Lead  monoxide  

PbO 

50, 

VI 


APPENDIX. 


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APPENDIX. 


Table  IV.  —  LIST  OF  THE   METALS  IN  THE  ORDER  OF 
DECREASING  SOLUTION  TENSION. 


The  Alkali  Metals. 

The  Alkaline-Earth  Metals. 

Magnesium. 

Aluminum. 

Manganese. 

Zinc. 

Cadmium. 

Iron. 

Cobalt. 

Nickel. 

Tin. 

Lead. 


Hydrogen. 

Bismuth. 

Antimony. 

Copper. 

Arsenic. 

Mercury. 

Silver. 

Palladium. 

Platinum. 

Gold. 


The  metals  appearing  first  in  the  table  can  replace  those  that 
follow,  in  the  solutions  of  their  salts. 


INDEX.* 


TNumbers  denote  pages,  or,  if  preceded  by  "  Ex.,"  Laboratory  Exercises  in 

Tart  II.] 


Acetic  acid 95 

Acetylene 207 

Acid  and  normal  salts,  Ex.  22,  Ex.  23 

properties 93 

radical 98 

salts 100 

salts,  converted  into   normal 

salts 100 

acetic     95 

carbolic 210 

carbonic 197 

chloric 266 

chlorous 266 

chlorplatinic 405 

lluosilicic 317 

hydriodic,    261,  262,  Ex.  44,  Ex.  45 

hydrobromic 259,  Ex.  45 

hydrochloric  .  88,  89,  90,  91,  Ex.  17 

hydrocyanic 204 

hydrofluoric 255 

hydrosulphuric     .  169, 170,  Ex.  33 

hypocblorous 265 

hyponitrous 163 

hypophosphorous 285 

iodic 268 

nitric  ...     95, 149,  Ex.  18,  Ex.  26 

nitric,  fuming 154 

nitrous 158 

perchloric 267 

periodic 268 

phosphoric  .  .  286,  287,  288,  Ex.  47 


Acid,  phosphorous     .......   286 

silicic 318 

stannic 404 

sulphuric,  95, 174 /.,  Ex.  18,  Ex.35 

sulphurous 183 

thiosulphuric 184 

Acidity  and  basicity 102 

Acids,  action  of,  on  bases  ....     93 

action  of,  on  metals 93 

action  with  oxides 99 

most  important  property  of   .     96 

nomenclature  of 103 

polysilicic 319 

properties  of Ex.18 

see  under  specific  names. 

strong  and  weak 330 

Agent,     drying,    for     hydrogen 

(calcium  chloride) 12 

drying,  for  hydrochloric  acid 

(sulphuric  acid) 89 

drying,  for  ammonia  (lime)  .    137 

Air,  a  physical  mixture 119 

determination   of  proportion 

of  oxygen 117 

liquefaction  of 115 

Alabaster 352 

Alkalies,  defined 97 

Alkali  metals 337 

distinction  between    .  .  .    Ex.  57 
properties  change  in  order  of 
atomic  masses 333 


*  This  Complete  Index  is  bound  in  the  book  with  and  without  the  Laboratory 
Exercises. 

ix 


INDEX. 


Alkaline  earth  metals 350 

Allot  ropism 167,  251,  272 

Alum,  crystallization  of    .   .    Ex.13 
water  of  crystallization  in,  Ex.  10 

Aluminates      380 

Aluminum  and  its  compounds, 

Ex.  69 

bronze 368 

carbonate  and  sulphide  .  .  .  382 
hydroxide  ....  Ex.69,d,  g 
occurrence  and  preparation 

of 378 

oxide  and  hydroxide 381 

properties  of 379 

salts 382 

uses  of  ...,,, 380 

Alums 382,  Ex.69,  e,/ 

Amalgam,  ammonium    .    348,  Ex.  57 
Amalgamation  process  for  silver  370 

Amalgams 365 

Ammonia,    action    of    chlorine 

on 85,  146 

a  refrigerating  agent 140 

chemical  properties  of  .  ...  142 
commercial  sources  of  ....  138 

composition  of 145 

existence  of 136 

from  illuminating  gas  manu- 
facture   209 

laboratory  methods  of  prepa- 
ration      136,  137 

liquefaction  of .   140 

physical  properties  of  ....  139 
preparation  and  properties 

of     Ex.  25 

process  of  making  soda  ...   341 

synthesis  of 145 

Ammonium     .  .  * 348 

amalgam Ex.  57 

compounds 142 

compounds,  dissociation  of  .   144 

hydroxide 97 

Amorphous  carbon 187 

artificial 188 

Analogues  of  an  element  .  .  .    309 
Analysis,  volumetric,  use  of  per- 
manganate in  394 


Anhydride  of  acid  defined   ...     96 

Aniline 157 

Animal     charcoal    and     bone- 
black     191 

Annealing  and  tempering  steel .    386 

glass 321,  Ex.2 

Antidote  for  arsenic 294 

Antimony .    Ex.  49 

black      300 

chemical  properties 298 

compounds 299 

physical  properties 297 

preparation 297 

trisulphide 300 

uses  of 300 

Aqua  regia 82 

a  source  of  nascent  chlorine  .   154 

Argon Ill 

family 314 

Arsenates 296 

Arsenic c  ,  .  .    Ex.  48 

acid 295 

chloride 291 

greens 295 

Marsh's  test  for 292 

occurrence    and    preparation 

of 290 

properties  of 290 

trioxide 293 

trisulphide 296 

Arsenious  acid 294 

double  nature  of 294 

Arsenite  of  sodium 294 

Arsine 292 

Asbestos 358 

Ash 188 

of  land  plants 201 

of  sea  plants 201 

Atomic  hypothesis 224 

masses,  determination  of,  233,  234 
masses,   exact,   obtained    by 
comparing  with  oxygen   .  .   235 

masses,  relative 232 

mass  methods,  application  of, 

237-240 

Atomic,  or  nascent,  state   .  .  242,243 
Atoms 224 


INDEX. 


XI 


Atoms   and    molecules,   distinc- 
tion between 226 

number  of,  in  molecules   of 
elements 243 

Atmosphere,  carbon  dioxide  in,  112 

character  of  the 109 

water  vapor  in 113 

weight  and  pressure  of     ...   114 

Atmospheric  dust 113 

Avogadro's  hypothesis    ....   128 

Baking:  powders 195 

Balanced,  or  equilibrium,  equa- 
tion      263,265 

Barium  and  strontium   ....   356 
Barium  chloride,  water  of  crys- 
tallization in Ex.  61 

hydroxide 97 

peroxide 356 

peroxide,  a  source  of  oxygen,    27 
Barometric  reading,  correction 

Of 126 

Bases,  nomenclature  of 105 

properties  of Ex.  19 

Basic  salts 101 

Basicity  and  acidity 102 

Bead,  borax Ex.  51 

Bell-metal 368 

Benzene      210 

Beryllium,  or  glucinum     ....   358 
Bessemer  converter,  Fig.  63    .  .   387 

process  for  steel 387 

Bicarbonates- 198 

Bismuth 300,  Ex.  50 

properties  of 301 

salts 301 

uses  of 302 

Blast-furnace,  Fig.  62 384 

Bleaching  powder 355 

with  chlorine     ...   86,  87,  Ex.  16 
with  sulphur  dioxide    ....   182 
Blowpipe,  oxidation  of  tin,  Ex.  75,  h 
oxidizing  and  reducing  flames, 

Ex.  63,  b 

oxyhydrogen Fig.  6,    14 

reduction  of  lead  oxide  .  Ex.  74,  6 
Blue  prints  ..,.,..,...  374 


Blue  vitriol,  crystallization  of,  Ex.  13 
water  of  crystallization  in,  Ex.  10 

Boiling-point  of  water 48 

Bone-black  and  animal  charcoal,  191 

Borax 323 

and  boric  acid Ex.51 

Borax  bead Ex.51 

Boric  acid 323 

Boron,  occurrence  and  prepara- 
tion     321 

properties  of 322 

Boyle's,  or  Mariotte's,  law    .   .    121 

Brass 368 

Bromates 268 

Bromine Ex.  43 

oxyhydrogen  compounds   .  .  267 

preparation  of 257 

properties  of 258 

Bronze 368 

Bunsen  burner 217,  Ex.  1 

flame 217,  Ex.  41 

Burner,  Bunsen Ex.  1 

Burning.    See  Combustion. 

Cadmium 363 

compounds Ex.  65 

Calcium 350 

carbide 207 

carbonate 352 

chloride 352 

chloride,  deliquescence  of,  Ex.  12 

compounds Ex.  58 

hydroxide 97,  351 

oxide 350 

phosphate 353 

silicate 356 

sulphate 352 

sulphide 356 

Calculation  of  quantities  of  fac- 
tors and  products 20 

Calomel 364 

Candle  flame,  Figs.  51  and  52, 

215,  216 

"  Carats  fine "    . 377 

Carbon 186,  Ex.  37 

amorphous 187 

amorphous,  artificial    ....   188 


Xll 


INDEX. 


Carbon  and  hyd-regen  compounds,  205 

coke  and  gas 190 

reduction  by Ex.  40 

Carbonates 197 

identification  of 198 

natural     200 

Carbon  dioxide  ...     Ex.  as,  Ex.  39 
chemical  properties  of  .  .  .  .   193 

in  the  atmosphere 112 

occurrence  of 191 

other  sources  of 194,  195 

ph}-sical  properties  ot  ....   192 

preparation 192 

relation  of,  to  life 196 

Carbon  disulphide 174 

Carbon  monoxide,  formation  of,  201 

laboratory  method 202 

properties  of 202 

reduces  iron  ore 3S5 

Carbonic  acid 197 

Carbonic  anhydride 197 

Carbonization,  natural;  coal  .   .    188 

Carborundum 318 

Cast-iron 385,386 

Caustic  potash 345 

Caustic  soda 340 

Celluloid 157 

Cement  and  mortar 355 

hydraulic     355 

Portland   . 355 

Chalk 352 

Chamber  crystals     .......    176 

Change,  chemical Ex.  3 

chemical,  and  energy  changes,  219 
Changes,  quantitative  cnaracter 

of  chemical 19 

Charcoal, animal,  and  bune-black,  191 

properties  of Ex.  37 

reduction  of  lead  oxme  by, 

Ex.  40,  a 

wood 189 

Charles' law 123 

Charring •....'  187 

Chemical  change Ex.  3 

fundamental  fact  of 66 

Chemical  changes  aim  er.ergy 

changes 219 


Chemical  changes,  quantitative 

character  of 19 

reaction  defined 66 

Chemistry  and  physics,  relation 

between    . 2 

definition  of s       1 

importance  of 1 

Chile  saltpeter 156,338,344 

Chlorates  and  chloric  acid    .   .    206 

Chlorides 93 

formed  by  dissolving  metals 

in  aqua  regia 82 

Chlorine,  action  of,  on   ammo- 
nia    85,  146 

and  hydrogen,  union  of   ...     84 

and  turpentine 86 

and  water 85 

chemical  properties  of  ....     84 
common  method  of  prepara- 
tion       79 

existence  of 79 

liquefaction  of 83 

other    methods    of    prepara- 
tion   81,  82 

preparation    and    properties 

of Ex.  16 

uses  of .' 86 

water,  preparation  of    Ex.  16,  c,  i 
Chlorine   dioxide  and    tetrox- 

ide 264 

Chlorine  hydrate 83 

Chlorine  monoxide 264 

Chlorous  acid 2(56 

Chlorplatinic  acid 405 

Chromates    and    dichromates, 

398,  399 
distinguished    from    chromic 

salts Ex.73,/ 

Chrome-alum 397 

Chromic  salts  changed  to  chrom- 

ites Ex.  73,  e 

distinguished  from  chromates, 

Ex.73,./ 

from  dichromates  or  chrom- 
ates   Ex.  73,  h 

Chromite    . 395 

Chromites     .  ...   397 


INDEX. 


Xlll 


Chromium 395 

compounds Ex.  73 

double  nature  of 397 

oxides  and  hydroxides  .   ...   390 
Chromous  and  chromic  salts,  396 

Clay 378,  3S3 

Coal,  natural  carbonization  ...   188 

"  Coal  gas  " 203 

Coals  and  wood,  composition  of,  189 

Cobalt 391 

Cobalt  and  nickel Ex.  71 

chloride  and  water  .  .  .    Ex.  71,/ 

Coin  gold 376 

Coin  silver 372 

Coke 210 

Coke  and  gas  carbon 190 

Collodion 157 

Couibining  proportions  ....     73 

use  of 74 

Combustion,    fixed    amount   of 

heat  of 31 

in  air 15 

in  air;  drafts 33 

in  chlorine 84 

in  oxygen .  Ex.  7 

oi'dinary 31 

reversed 36 

slow 31 

spontaneous 32 

supporting 15 

temperature  depends  on  rate 

of 31 

Commercial  iron 386 

Comparison    of   the    halogen 

acids Ex.  45 

Compound,  defined  .......       5 

of  two  elements,  how  named,    74 
Compounds,  molecules  of     ...    227 
of    same  two  elements,  how 

distinguished 75,  77 

Conductivity,  electric,  of  solu- 
tions     326 

Constant  proportions    .   .   .    Ex.  15 

law  of 67 

Converter,  Bessemer,  Fig.  63   .   .    387 

Copper 366 

alloys 368 


Copper  and  its  compounds  .    Ex.  67 

compounds 368 

etc.,  relation  to  alkali  metals  .   366 

native .'....   366 

ores 307 

properties  and  uses 367 

sulphate  use  in  making  hydro- 
gen       11 

Copper-plating      369 

Coral 201 

Correction   of  the  barometric 

reading 126 

Corrosive  sublimate 365 

Critical  temperature  defined     .    116 
Crucible  process  for  steel  .   .   .    387 

Cryolite 378 

Crystallization 63,  Ex.  13 

from  fusion 64 

water  of Ex.10 

Crystal-water  in  barium  chlor- 
ide  . Ex.61 

in  gypsum Ex.59 

Cupric  oxide 369,  Ex.  67,  c 

Cupric  sulphate 3(59 

Cupric  sulphide 369 

Cuprous  oxide    ....  368,  Ex.  67,  e 
Cyanide  process  of  extracting 

gold 375 

Cyanogen 203 


Davy's  safety  lamp 35 

Deacon's  process  for  chlorine  .     82 
Decantation  and  filtration    .  .     64 

Decay  of  wood 31 

Decomposition,  heat  of    ....    220 
Decrepitation     54,  Ex.  23 

of  salt 344 

Definite    proportions,    law    of, 

explained  by  atomic  theory,  224 
Deflagration 30 

of  potassium  nitrate  and  sul- 
phur   Ex.  55 

Dehydrating  or  drying  agents, 

12,  56,  89,  137,  346 

Deliquescence   ....  56, 131,  Ex.  12 
Density,  methods  of  vapor  .  .  .   228 


XIV 


INDEX. 


Developing,  photographic   ...    373 

Dew-point 113 

Diamond 186 

Diatoms     315 

Dichromates  and  chromates, 

398,  390 
Diffusion,  explanation  of  .   ...    129 

of  hydrogen 17 

Dioxides  and  peroxides  ....    277 
Disodium  hydrogen  phosphate 

and  magnesium  salts,  Kx .  62,  d 

Dissociation  by  heat 325 

iu  solution,  ionization   ....   325 
of  ammonium  compounds  .  .    144 

of  arsenic  molecules 291 

of  hydi-iodic  acid 262 

of  hydrobromic  acid 260 

of  iodine 261 

of  nitrogen  tetroxide 160 

of  steam 48 

Distillation,  apparatus  for  ...     46 

defined     45 

Distilling  at  reduced  pressure,  275 

Dolomite 201 

Double  cyanide  of  silver  and 

potassium Ex.68,/ 

Drafts,  combustion  in 34 

Drying  agent 56 

calcium  chloride  .  .  .    12,137,352 

calcium  oxide 56 

potassium  carbonate 346 

sodium  hydroxide  ......   137 

sulphuric  acid 89 

Dulong  and  Petit's  rule  ....    235 

Dust,  atmospheric 113 

Dynamite 157 

Earthenware 383 

Effervescence 04 

effect  of  pressure  on  ...     Ex.  39 

Efflorescence 55, 131,  Ex.  11 

Electric  conductivity  of  solu- 
tions   326 

Electric  current,  to  extract  met- 
als from  compounds  ....   335 
Electric  furnace,  Figs.  48  and  61, 

207,  379 


Electrolysis 329 

of  hydrochloric  acid,  Fig.  23, 

83,  91 

of  sodium  chloride 83 

of  water 19 

preparation  of  aluminum  by,  370 
preparation  of  potassium  by  .    345 
preparation  of  sodium  by  .  .   339 
preparation    of    sodium    hy- 
droxide by 340 

Electrolytes      19, 327 

Electrotype  plates 369 

Element,  analogues  of  an    ...   309 
Elements,  abundance  of   ....       7 

defined 5 

heterologous 300 

homologous 305 

importance  of 7 

list  of     6,  Appendix 

molecules  of 227 

number  of  atoms  in  molecule 

of 24.J 

periodic  table  of 312,313 

prediction  of  unknown    .  .  .    311 
properties  of,  determined   .  .   309 
Energy      changes     accompany 

chemical  changes 219 

potential,  of  elements    ....  219 
Equation,   equilibrium    or    bal- 
anced    263,265 

ionic 3_'S 

volumetric  meaning  of  .  .  .  .   246 
Equations  and  molecular  formu- 
las     241 

and      symbols,     quantitative 

meaning  of 70 

symbolic 68 

the  result  of  experiment .  .  .     70 
Equilibrium  equation  of  a  solu- 
tion      327 

in  solutions  of  gases  and  sol- 
ids    131,132 

or  balanced  equation     .  .  263,  265 
unstable,  of  compounds  with 
negative  heat  of  formation,  221 

Equivalent 237 

of  magnesium  .......  Ex.  (I 


INDEX. 


XV 


Equivalent  of  zinc   .   .   .    237,  Ex.  64 

Etching  of  glass 256,321 

Ethane 206 

Ethylene 206 

Eudiometer 40 

Evaporation 131,  Ex.  4 

Explosion  of  hydrogen,  velocity 

of     16 

Factor  denned 4 

Factors  and  products     .  .    20,  69,  80 

Family,  argon  .   .  .   „ 314 

calcium 359 

halogen 268 

nitrogen 303 

Families,  natural 305 

Feldspar 378 

Fermentation 194 

Ferric  chloride 390 

Ferric  hydroxide     389 

Ferric  oxide 389 

Ferric  sulphate 390 

Ferrous  ammonium  sulphate,  390 

Ferrous  chloride 389 

Ferrous-ferric  oxide  ......    389 

Ferrous  hydroxide  .......    389 

Ferrous  sulphate 390 

Fertilizers 354 

Filtrate 64 

Filtration Ex.  4 

Fire-damp 205 

Fixing,  photographic 374 

Flame,  Bunsep,  Fig.  54 217 

candle,  Figs.  51  and  52  .  .  215,  216 
colored  by  barium  .  .  357,  Ex.  60 
colored  by  calcium  .  .  .  .  Ex.58 
colored  by  potassium,  345,  Ex.  55 
colored  by  sodium  .  .  344,  Ex.54 
colored  by  strontium  .  357,  Ex.  CO 

defined 35,  214 

hydrogen,  Fig.  7 14 

non-luminous 216 

oxidizing,  with  blowpipe,  Ex.  63, 6 
reducing,  with  blowpipe,  Ex.  63,  6 

Flames Ex.  41 

dissection  of,  Fig.  55 218 

luminosity  of ,   214 


Flames,    oxidizing    and    reduc- 
ing   Ex.51 

simple  and  complex 217 

structure  of 215 

Flask,  generating 10 

Fluorine 254,255 

Fluosilicic  acid 317 

Flux 254 

used  in  blast-furnace    ....  384 

Fool's  gold 389 

Formula  types  based  on  valence,  248 

Formulas  and  symbols 68 

and  symbols,  how  to  rcpi'escnt 

multiples  of 71 

graphic,  or  structural    .  .  249,  25J 

how  determined 240 

molecular,  and  equations    .  .  241 
Freezing-point  of  water    ....     48 

Fuming  nitric  acid     15t 

Furnace,  blast,  Fii,'.  62 384 

electric,  Figs.  48  and  61  .  207,  379 
Hall's  aluminum,  Fi^.  61  .  .  379 
reverberatory,  Fig.  60  ....  371 

Gallium,  properties  predicted    .   311 

Galvanized  iron 362 

"  Gaps  "  in  the  periodic  arrange- 
ment   310 

Gas,  amount  used 213 

carbon 210 

carbon  and  coke 190 

collection  of,  over  water, 

10,  Ex.  5 

collection  of,  by  upward  dis- 
placement    11 

collection   of,   by  downward 

displacement 80 

comparison  of  the  two  kinds 

of  illuminating 212 

permanent 116 

relation  of  pressure  to  volume 

of 121 

relation  of  temperature  to  vol- 
ume of  122 

Gaseous  substances,  molecular 

masses  of 227 

Gases  and  vapors  defined  ....   121 


XVI 


INDEX. 


Gases,  diffusion  of 120,  130 

solution  of,  in  liquids    ....   131 

Generating  flask 10 

Generator,  Kipp's,  Fig.  4  ....     12 
Germanium,     properties     pre- 
dicted      311 

German  silver 36S 

Glass 320 

color  of     320 

cut 321 

etching  of 250,  321 

prcs^el 321 

tubing,  cutting  and  bending,  Ex.  2 

Glauber's  salt 344 

effloresce  nee  of Ex.  11 

Gluciimtn 358 

Glycerine 340 

Gold 375 

fineness,  in  carats 377 

properties  and  uses  of  ....   376 

Granite 378 

"  Granite  ironware  " 320 

Grape-sugar     and     potassium 

permanganate     .   .   .  Ex.72,  d 
Graphic,  or  structural,  formulas, 

249,  250 

Graphite 187,  Ex.  37 

retorts  used  in  distilling  zinc,  371 

Greens,  arsenic 295 

Green  vitriol 390 

Group,  calcium 359 

zinc .   361 

Gun  cotton 157 

metal 368 

powder 157 

Gypsurn 352 

water  of  crystallization  in,   Ex.  59 

Haematite 384 

Hall's  process  for  making  alu- 
minum      378 

Halogens 254 

Halogen  acids,  comparison  of,  Ex.  45 

family 268 

Halogen-oxygen  compounds  .    263 
Halogen     oxyhydrogen     conn- 
pound* 264,265 


Hardness,  permanent 47 

temporary 47 

Heat,  dissociation  by 325 

of  formation  and  decomposi- 
tion   220,221 

of  formation  evolved  in  stages,  221 
of    formation,    positive    and 

negative -221 

Heating  liquids  in  beakers,  etc., 

Ex.  4 
with     Bunsen      flame,     best 

method  qf Ex.  1 

Helium Ill 

Heterologous  elements    ....    309 
Homologous  elements     ....    305 

Hydraulic  cement 35.") 

mining  of  gold 375 

Hydriodic  acid  .    261,  Ex.  44,  Ex.  45 
powerf ul  reducing  agent  .  .  .   2G3 

properties  of 262 

Hydrobromic  acid,  .    .  .    259,  Ex.  45 
Hydrochloric  acid,  commercial 

method  of  preparation  .  *s>     89 
common  method   of  prepara- 
tion          88 

electrolysis  of 91 

existence  of 88 

physical  properties 90 

preparation  and  properties,  Ex.  17 

synthesis  of 91 

volumeti  ic  composition  of  .  .     91 

Hydrocyanic  acid 204 

Hydrofluoric  acid 255 

Hydrogen    and    chlorine,  union 

of 84 

antimonide 29!) 

rsenide 292 

methods  of  preparation   ...     I'.) 

occlusion  of 16 

peroxide 274,  Ex.  46 

peroxide,   composition  of  .  .   276 

peroxide,  test  for 276 

phosphide,  Fig.  58 282 

physical  properties  of  ....     17 
preparation     and    properties 

of     Ex.  5 

reducing  power  of  ......    10 


INDEX. 


XV11 


Hydrogen  silicule 316 

sulphide 169,  Ex.33 

sulphide,  properties  of  ....   170 

velocity  of  explosion  of  ...     16 

Hydrogen-carbon  compounds,  205 

Hydrolysis 330,331 

and  mass  action Ex.53 

of    aluminum    carbonate  and 

sulphide 382 

of  magnesium  carbonate    .  .   358 

of  soap     341 

Hydrosulphides 172 

Hydrosulphuric  acid 1G9 

Hydroxides 52 

action  of  metals  upon   ....     52 

Hydroxyl 98 

Hydroxylamine »    .   .    152 

Hygroscopic 56 

"Hypo" 184 

Hypobromites 267 

Hypochlorites 265 

Hypochlorous  acid 265 

Hyponitrous  acid 163 

Hypophosphorous  acid  ....    285 

Hypothesis,  Avogadro's   ....    128 

atomic 224 

Iceland  spar 201 

Ignition  temperature 32 

Illuminating  gas 208 

comparison  of  the  two  kinds  .  212 
distillation  of  coal,  Fig.  49  .  .  208 
water  gas  process,  Fig.  50  .  .  210 

Ink,  sympathetic 391 

lodates 268 

lodicacid 268 

Iodine Ex.  44 

occurrence  and  preparation  .   260 
Iodine     oxyhydrogen    com- 
pounds     268 

Iodine  pentoxide 264 

properties  of 261 

Ionic  equation 328 

lonization 325,  327,  Ex.  52 

Ions .T.  7  .   .    327 

not  atoms 330 

Iron Ex.  70 


Iron  chlorides 389 

commercial 386 

compounds,  comparison  of  the 

two  classes Ex.  70,  e 

galvanized 362 

occurrence  and  metallurgy    .    384 
oxides  and  hydroxides  ....   388 

properties  of 388 

pyrites 169,  384,  389 

sulphates 3!)0 

sulphides     389 

Isouierisml 250 

Kaolin 319,383 

Kindling  temperature     .    32,  Ex.  8 
Krypton 314 

Lampblack 189 

Latent  lieat  of  water 48 

Law,  TSoyle's  or  Mariotte's  .  121,122 

Charles' 123 

of  conservation  of  matter  .  .     67 
of   constant    proportions    by 

weight 67 

of  definite  proportions  ex- 
plained by  the  atomic  the- 
ory   224 

of  multiple  proportions  .  .  .   223 
of  multiple  proportions,  ex- 
planation of 225 

of  persistence  of  mass  ....     67 
Laws  of  osmotic  pressure     .  .  .   133 
of   simple  and   multiple  vol- 
umes   215 

Lead  and  its  compounds     .  .    Ex.  74 
occurrence,  preparation,  and 

properties 400 

oxides,  nitrate  and  acetate  .  .  401 
sulphate,   chloride,  and   car- 

bonate 402 

sulphide 402 

uses  of 401 

Le  Blanc  process 89,  341 

Life,  relation  of  carbon  dioxide 

to     196 

Light,    action    on    silver    com- 
pounds   373,  Ex.  68 


XV111 


INDEX. 


Lime 350 

"air  slaked" 351 

slaking     Ex.58 

Limestone 200,  352 

Liquefaction  of  ammonia  .  .  .  140 
Liquid  air,  properties  of  ....  117 
Liquids,  diffusion  of  ....  129,130 

Lithium      338 

Litmus  solution,  effect  of  char- 
coal on Ex.  37 

Lodegtone     389 

Luminous  flames Ex.  41 

Luster,  metallic 334 

Magnalium 381 

Magnesite     201 

effect  of  heat  upon,  Ex.  40,  Ex.  62,  e 

Magnetite 384,  389 

Magnesium  . 357 

ammonium  phosphate  .  .    Ex.47 
and  its  compounds  .   .  .  .    Ex.62 

carbonate 358 

chloride 358 

equivalent  of Ex.6 

oxide 357 

silicide 317 

sulphate 358 

Manganatea 393 

Manganese 392 

compounds Ex.  72 

dioxide   and   hydrogen   per- 
oxide   Ex.  46 

dioxide  a  source  of  oxygen   .     27 
dioxide,  action  on  potassium 

chlorate 26 

dioxide,  to  liberate  chlorine, 

Ex.16 
dioxide,  to  prepare  oxygen,  Ex.  7 

oxides 392 

Manganites 395 

Marble 201,352 

result  of  heating Ex.58 

Mariotte's,  or  Boyle'g,  law,  121,122 

Marl 352 

Marsh  gas 205 

•  Marsh's  test  for  antimony    ...   299 
for  arsenic 292 


Mass,  persistence  of 66 

Mass  action 331,332 

and  hydrolysis Ex.53 

Matches 281 

Matter,  law  of  conservation  of  -     67 

physical  states  of 123 

Mercuric  chloride 365 

oxide 364 

oxide,  decomposition  of,  27,  Ex.  3 

sulphide 365 

Mercurous  chloride 364 

oxide 364 

Mercury 363 

and  its  compounds  ....    Ex.  66 

compounds 364 

Metal,  Rose's 302 

type 297 

Wood's 302 

Metallic  luster 334 

Metalloids 331 

Metals,  alkali 337 

alkali,  distinction  between,  Ex.  57 

alkaline-earth 350 

and  non-metals 334 

extraction  of,  from  ores  .  335,  385 

occurrence  of 334 

properties  of 335,  336 

"  Metaphosphate  head  "     .  '.  .    288 

Methane 205 

Mica 378 

Mineralg 334 

Mirrors 373 

Molecular   formulas  and   equa- 
tions   ...       241 

mass,  boiling-point  and  freez- 
ing-point methods 230 

mass,  exact  methods  of  obtain- 
ing    231 

mass  of   oxygen,  32,   reason 

for .   244 

mass,  other  methods  of  deter- 
mining   229 

mass,  oxygen  the  standard  of,  228 
mass,  vapor  density  methods 

for 228,2-29 

masses  of  gaseous  substances,  227 
theory 127 


INDEX. 


XIX 


Molecules  and  atoms,  distinction 

between 226 

of  elements  and  of  compounds  227 
of  elements,  number  of  atoms 

in 243 

Mordant,  aluminum  compounds,  382 
Mortar,  action  of  acid  on  old,  Ex.  58 

anil  cement 355 

Multiple  proportions,  explana- 
tion of  law  of 225 

law  of 223,225 

Multiples  of   symbols  and  for- 
mulas, how  represented  ...     71 

Naming  a  compound  of  two  ele- 
ments          74 

different  compounds  of  same 

two  elements 75-77 

Naphthalene 210 

Nascent,  or  atomic,  state    .   .  242,243 

Natural  families 305 

Natural  family  of  elements    .  .   270 

Neon 314 

Neutralization 98,  Ex.  21 

explanation  of 327, 328 

Neutralize 96 

Nickel 391 

and  cobalt Ex.71 

Nitrate,  manufacture  of   potas- 
sium   ..." 156 

Nitrate  beds 156 

Nitrates  and  nitric  acid,  uses  of  157 
Nitrates, formation  of,  in  nature,  155 

Nitre  .   .   .    .  T 346 

Nitric  acid 95,149 

action  on  metals 152 

and  nitrates,  uses  of 157 

commercial  preparation  .  .  .   149 
explanation  of  oxidation  by  .   154 

faming b 154 

laboratory  preparation  of  .  .   150 

oxidation  by 154 

powerful  oxidizing  agent   .  .   152 

preparation Ex.  26 

preparation    compared    with 

that  of  hydrochloric  acid  .  .    150 
properties  of  .  .  151,  Ex.  18,  Ex.  26 


Nitric  acid,  oxidizing  agent, 

Ex.  34,  e;  Ex.  70,  d;  Ex.  75,  9 
reduced  by  nascent  hydrogen,  152 

Nitric  anhydride 158 

Nitric  oxide 161,  Ex.29 

Nitrides      109 

Nitrites 159,  Ex.  27 

Nitrobenzene 157 

Nitrogen  dioxide  and  tetroxide, 

160,  Ex.  28 

existence  of    .  .  .  , 107 

family 303,304 

from  ammonium  nitrite    .  .  .   108 

of  the  air,  a  mixture 109 

oxides 155 

pentoxide 158 

preparation  and  properties, 

E*.  24 
preparation  of,  from  air,  Figs. 

25  and  26 107 

properties  of 108 

trioxide 109 

Nitrous  acid 158 

anhydride 158 

oxide 102,  Ex.  30 

Nomenclature  of  acids  ..   ...    103 

of  bases 105 

of  salts 104 

Non-luminous  flames    .  21(5,  Ex.41 

Normal  salts 99 

and  acid  salts    .  .    Ex.  22,  Ex.  23 

Occlusion  of  hydrogen  ....  16 
Open  hearth  process  for  steel,  387 
Ores 334 

extraction  of  metals  from  .  .  335 

Osmotic  pressure     .  • 132 

laws  of 133 

Oxidation  by  chromates  and  di- 

chromates 399 

by  nitric  acid     .  .  .   154,  Ex.  75,  e 
by  nitric  acid,  explanation  of,  154 

by  permanganate 394 

and  reduction 30 

Oxides      29 

of  phosphorus 284 

with  acids,  action  of  .....     99 


A 


INDEX. 


Oxidizing  agent    ........     30 

bromine    ...........   258 

chlorine    ........  Ex.  73,  g 

chromates  and  dichromates   .   399 
hydrogen  peroxide  ......   275 

nitric  acid    .....  154,  Ex.75,  e 

ozone  .............   273 

permanganate  .....  ...  394 

Oxidizing  flame    ......     Ex.51 

Oxygen  acids  of  phosphorus   .   .    285 
chemical  properties  of     ...     28 
determination  of  the  propor- 
tion  of,  in  air,  Figs.  30  and 
31  ............  117,118 

physical  properties  of  ....     27 

preparation  and  properties,  Ex.  7 
preparation  of  ....  24,  25,  26,  27 

reason  for  choosing  molecular 
mass  of,  32  ...  ......   244 

standard  of  atomic  mass  ...   232 
standard  of  molecular  mass  .   228 
weight  of  one  liter  ....    Ex.  42 

Ozone    ..............    273 

Palladium,  occlusion  of  hydro- 

gen  by    ...........     17 

Paris  green  .........  .  .  295 

Parke's  process  of  desilverizing 

lead    ............    371 

Peat  .......  .  .......   188 

Perchlorates    ..........    267 

Perchloric  acid  .....  ....    267 

Periodic  acid  ..........   268 

Periodic  arrangement    .   .   .  306JT. 
conclusion  ..........    312 

gaps  in  ............   310 

regularitie»in    ........  308 

Periodic  law    ..........   308 

Permanent  gases  ......  *.  .    116 

Permanganate,  oxidation  by     .   394 
Peroxides  and  dioxides  ....   277 

Peroxide  of  hydrogen  .    .   .     Ex.  46 
Phenol    .............  210 

Phenomena,  chemical  and  phys- 

ical  .............      3 

Phosphates,  necessary  for  plants,  354 
uses  of  .  ...........  289 


Phosphide  of    hydrogen,  Fig. 
58  ...............   282 

Phosphides  ...........    284 

Phosphine,  Fig.  58  .......   282 

Phosphonium  salts   ......    283 

Phosphorescence     .......    168 

Phosphoric  acids    .......   28(i 

preparation  of   ........  287 

salts  of  ............    288 

Phosphorous  acid  .......    286 

Phosphorus     ........  ,  .   279 

and  phosphoric  acid  .  .  .    Ex.47 
molecular  mass  of  ......  281 

oxides   ............   284 

properties  of  .........   279 

red  ..............  280 

Photography  ........   .   .    373 

Physical  phenomena    .....       3 

Physics,  relation  between  Chem- 
istry  and  ...........      2 

Pig-iron  ......  ,  ......   38;i 

Placer-mining  of  gold    .....   375 

Plaster  of  Paris    ........   352 

making  casts  with  ....    Ex.  58 

Platinum,  occlusion  of  hydrogen 

by    ............  16,  17 

occurrence  and  preparation  .   404 
properties  and  uses   .....  40.1 

Plumbites  and  plumbates     .   .   402 
Pneumatic  trough  !......     10 

Polysilicic  acids    ........    319 

Porcelain-lined  ware     .....   320 

Porcelain,  stoneware,  etc.      .   .   383 
Portland  cement  ...*....    355 

Potash     ...........  346,  347 

Potassium     .........  ,  .   345 

bromide    ...........  347 

carbonate    ........  201,  340 

chlorate    ...........  346 

chloride,  solubility  of  .  .    Ex.  5<; 
compounds  ........    Ex.  55 

ferricyanide  .........  391 

ferrocyanide      *  .......  390 

hydrogen  tartrate  .  Ex.  55,  Ex.  57 
hydroxide  .........  97,  345 

hydroxide,  deliquescence  of, 

Ex.  12 


INDEX. 


XXI 


Potassium  iodide 347 

nitrate 346 

nitrate,  manufacture  of, 

156,  Ex.  55 

permanganate 393 

silicate 319 

Powders,  baking 195 

Precipitant 63 

Precipitate 63 

Precipitation Ex.  14 

and  crystallization 63 

Prediction    of    unknown    ele- 
ments     311 

Pressure   and   temperature,  re- 
duction to  standard    ....   124 

osmotic 132 

relation  of,  to  volume  of  a  gas,  121 

Products  and  factors 69 

calculation  of    quantities  of 

factors  and 20 

Properties,  chemical  and  physi- 
cal        4 

effect  of  structure  on     ....   277 

of  acids Ex.  18 

of  bases Ex.  19 

of  elements  determined    .   .  .  309 

of  salts Ex.  20 

Proportions,  combining    ....     73 

constant 67,  Ex.  15 

law  of  multiple 223,  225 

use  of  combining 73 

Prussian  blue    ....  391,  Ex.  70,  d 

Pyrites 174,  389 

Pyrolusite  .  _." 27,  392 

Quantitative  meaning  of  sym- 
bols and  equations  ....     70 
Quicklime 350 

Radical,  acid 98 

metallic 143 

Reaction  defined 4,  66 

Reactions,  classification  of  ...     66 

Reagent  defined 4 

Red  phosphorus 280 

precipitate 364,  Ex.  3 

Reducing  agent,  hydriodic  acid,  263 
hydrogen  peroxide 276 


Reducing  agent,  nascent  hydro- 
gen   152 

Reducing  flame     .   .     Ex.  51,  Ex.  63 

Reduction 30 

by  carbon Ex.  40 

by  hydrogen 42 

of  potassium  nitrate  by  lead, 

Ex.  27 
Refrigerating   agent,  ammonia 

as  a     140 

Relation  of  volume  of  a  gas  to 

pressure     121 

Reverberatory  furnace    ....   371 

"  Reversion  to  type  " 307 

Rose's  metal 302 

Safety  lamp 35 

.matches 281 

Salt 343 

acid 100 

defined      99 

"Epsom" 358 

normal 99 

Saltpeter 346 

Salts,  basio ,  .   101 

nomenclature  of 104 

normal  and  acid  .     Ex.  22,  Ex.  23 

properties  of Ex.  20 

Saponify 340 

Saturated  solution 62 

Scandium,  properties  predicted,  311 

Scheele's  green 295 

Schweinfurt's  green 295 

Sea-water .     45 

Semipermeable  cell 133 

Shells 201 

Siderite,  or  spathic  iro«  ore     .  .   384 

Silicates 319 

Silicic  aoid 318 

anhydride ,  ....  318 

Silicon  carbide ,   .  .   318 

compounds 316 

dioxide 318 

hydride     316 

occurrence  and  preparation  .   315 

tetrachloride 317 

tetrafluoride  .  ....   317 


XX11 


INDEX. 


Silver 370 

and  its  compounds     ....  Ex.  68 

bromide    .  .  . 373 

chloride 373 

"  coin  " 372 

compounds  .  .  . 373 

extraction  of 370 

iodide 373 

nitrate 373 

-plating     372 

properties  and  uses *372 

"sterling" 372 

Slag 254 

Slaking  of  lime 351,  Ex.  58 

Smelting  of  silver 370 

Soap 340 

hydrolyzed  In  solution  ....   341 

Soda-ash 342 

Soda,  crystallized 342 

Sodium 338 

action  upon  water  .   .  .   49,  Ex.  9 

amalgam 92, 339 

and  water,  quantitative  study 

of  the  reaction 52 

arsenite 294,  Ex.  48 

bicarbonate  .  .  343,  Ex.  39,  Ex.  54 

carbonate 201,  341,  Ex.11 

chloride 343 

compounds Ex.54 

dichromate 399 

hydroxide 90,340 

nitrate 156,344 

oxides 339 

phosphate 343 

preparation  and  properties, 

49,  339,  Ex.  9 

sulphate 344,  Ex.  11 

sulphite Ex.34 

zincate Ex.  63,  d,  e 

"  Softening  "  water  with  goda,  353 

Solids,  diffusion  of 129 

solution  in  liquids 131 

Solubility  ...... 60 

of  potassium  chloride  .  .    Ex.  56 
Soluble  and  insoluble  substan- 
ces        61 

Solution     .  . £x.  4 


Solution,  boiling  point  of  ....     59 
character  of    .........     58 

composition  of  a 327 

effect  of  temperature  on,  60,  Ex.  13 
equilibrium  equation  of   ...   327 

freezing-point  of 59 

of  gases  in  liquids 131 

of  solids  in  liquids 131 

of  starch Ex.44 

saturated     62 

supersaturated     62 

temperature  changes  during  ,     59 
Solutions,  electric   conductivity 

of 326 

Solvay,  or  ammonia,  process  .  .   341 

Solvent 58 

Specific  heat  defined   .  •(  .  .   .   .    285 

of  water 47 

table Appendix 

Spelter 301 

Stalactites,  formation  of,  Fig.  46,  200 
Standard  temperature  and 

pressure,  reduction  to  .   .  .   124 

Stannic  acid     404 

and  staunous  compounds, 

404,  Ex.  75 

Starch  solution     Ex.44 

States  of  matter,  physical  .    .   .    128 
Steam  and  its  dissociation     ...     48 
volumetric  composition  of  .  .     41 
Steel,    properties  and  manufac- 
ture of   3S6,  387 

Sterling  silver 372 

Stibine 299 

Stoneware,  porcelain,  etc.  .  .  383 
Strontium  and  barium,  356,  Ex.  60 
Structural,  or  graphic,  formulas, 

249,  250 
Structure,  effect  of,  on  properties,  277 

Sublimation     . 144 

Substance,  compound    and    ele- 
mentary   .  .   . 5 

Sulphantimonites 300 

Sulpharsenious  acid 296 

Sulphates      179,  Ex.  36 

Sulphides,  formation  of    ....   171 
precipitation  of 172 


INDEX. 


XX111 


Sulphites 183 

Sulphostannates 404 

Sulphur,  chemical  properties  of, 

167,  Ex.  32 

compounds  of 1(>9 

dioxide     181,  Ex.  34 

dioxide,  properties  of  ....  182 
occurrence  and  preparation  .  165 
physical  properties  of, 

165, 166,  Ex.  31 

trioxide 180 

uses  of 168 

water 44 

Sulphuric  acid 95 

and  sulphates,  test  for  ....   180 

diluting 88,  89,  Ex.  17 

hydrates  of 177 

manufacture  of 174 

properties  of .    177,  Ex.  18,  Ex,  35 

purification  of  . 176 

reduction  of 177 

uses  of 179 

Sulphuric  anhydride 180 

Sulphurous  acid   . 183 

Sulphurous  anhydride    ....    183 
Supersaturated  solutions   ...     62 

Suspension Ex.  4 

Symhol  defined 5 

Symbolic  equations 68 

Symbols,  advantages  of  use  of  .     69 
and    equations,    quantitative 

meaning  of 70 

and  formulas 68 

and  formulas,  how  to  repre- 
sent multiples  of 71 

Sympathetic  ink 391 

System,  periodic 300  JT. 

Table  of  atomic  masses  .    Appendix 
of  specific  heats  .  .  .    Appendix 

the  periodic 312, 313 

of  metric  equivalents,  etc., 

Appendix 
of  vapor  tension  of  water, 

Appendix 

Tartar  emetic 300,  Ex.  49 

Tart  aric  acid 96 


Temperature  and  pressure,  re- 
duction to  standard    ....   124 

critical 116 

effect  of,  upon  solution,  60,  Ex.  13 

kindling 32 

of  ignition 32 

relation  of  volume  of  a  gas  to,  122 

.  Tern pering  and  annealing  steel,  386 

"  Test "  reactions  are  ionic  .  .  .   328 

Theory,  atomic 224 

ionization 325 

molecular 127 

Thiosulphates 184 

Tin  and  its  compounds    .    403,  Ex.  75 
occurrence  and   preparation 

of        402 

properties  and  uses 403 

Tincture Ex.  44 

Toluene 210 

Toning,  photographic 374 

Transition  elements  ....  295,  299 

Transpiration     .   .  < 18 

Trough,  pneumatic 10 

Tube,  safety,  or  thistle 10 

Tubing,    cutting    and     bending 

glass Ex.  2 

Turnbull's  blue Ex.  70 

Type,  reversion  to 307 

Type-metal 297 

Ultramarine 383 

Valence 241 

different  formula  types  based 

on 248 

of  members  of  argon  family,  314 
Vapor  density  methods  .  .  228,  229 
Vapors  and  gases  defined  ...  121 

Vein-mining  of  gold 375 

Victor  Meyer's  method  for  va- 
por density     229 

Volume  of  a  gas,  relation  of  pres 

sure  to 121 

of  a  gas,  relation  of  tempera- 

tureto 122 

Volumes,  laws  of    simple  and 

multiple 245 


XXIV 


INDEX. 


Volumetric  analysis,  use  of  per- 
manganate   394 

composition  of  steam    ....     39 
composition  of    hydrochloric 

acid 91 

composition  of  ammonia,   145,146 
meaning  of  an  equation   ...   246 

Water,  action  of  sodium  upon, 

49,  50,  Ex.9 
and  chlorine,  action  of  ....     85 

and  potassium 51 

as  integral  part  of  a  substance,    55 

distilled 47 

electrolysis  of    .......  19, 38 

formation  of 15 

hard  and  soft 47 

in  combination     53 

mechanically  enclosed  ....     53 
natural,  and  its  impurities  .  .     43 

nature  of .  . 38 

of  crystallization     ...  54,  Ex.  10 
of  crystallization  in  barium 

chloride Ex.  61 

of  crystallization  in  gypsum, 

Ex.59 

properties  of 47 

purification  of 45 


Water,  sea ,„....  45 

"  softened  "  with  soda  ....  353 

sulphur 44 

synthesis  of,  by  weight    .  .  .  42 

synthesis  of,  by  volume  ...  39 

vapor  in  the  atmosphere  ...  113 

-gas  process 210 

-glass 319 

Weight,     constant    proportions 

by 67 

of  one  liter  of  oxygen  .  .    Ex.  42 

White-lead 402 

Wood  charcoal 189 

Wood's  metal 302 

Wrought-iron 386 


Xenon 


314 


Zinc 361,  Ex.63 

chloride 363 

compounds 362,  Ex.  63 

dust 361 

equivalent  of Ex.  64 

group 361 

oxide 362 

sulphate  and  sulphide  ....  363 
used  in  desilverizing  lead  .  .  371 
uses  of 362 


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